Chapter 6 Thermochemistry 許富銀

Transcription

1 Chapter 6 Thermochemistry 許富銀

2 6.1 Chemical Hand Warmers Thermochemistry: the study of the relationships between chemistry and energy Hand warmers use the oxidation of iron as the exothermic reaction:

3 Nature of Energy Even though chemistry is the study of matter, energy affects matter. Energy is anything that has the capacity to do work. Work is a force acting over a distance. Energy = work = force distance Heat is the flow of energy caused by a difference in temperature. Energy can be exchanged between objects through contact. For example, through collisions

4 Classification of Energy Kinetic energy is energy of motion or energy that is being transferred. Thermal energy is the energy associated with temperature. Thermal energy is a form of kinetic energy.

5 Classification of Energy Potential energy is energy that is stored in an object, or energy associated with the composition and position of the object. Energy stored in the structure of a compound is potential energy.

6 FIGURE 6.1 The Different Manifestations of Energy

7 Conservation of Energy The law of conservation of energy states that energy cannot be created nor destroyed. When energy is transferred between objects, or converted from one form to another, the total amount of energy present at the beginning must be present at the end.

8 System and Surroundings We define the system as the material or process within which we are studying the energy changes within We define the surroundings as everything else with which the system can exchange energy.

9 Comparing the Amount of Energy in the System and Surroundings during Transfer Conservation of energy means that the amount of energy gained or lost by the system has to be equal to the amount of energy lost or gained by the surroundings.

10 Units of Energy The amount of kinetic energy an object has is directly proportional to its mass and velocity. The SI unit of energy is therefore kg m 2 /s 2, defined as the joule (J),

11 Units of Energy A calorie (cal) is the amount of energy needed to raise the temperature of one gram of water 1 C. kcal = energy needed to raise 1000 g of water 1 C food Calories = kcals

12 The First Law of Thermodynamics: Law of Conservation of Energy The first law of thermodynamics is the law of conservation of energy. This means that the total amount of energy in the universe is constant. According to the first law, a device that would continually produce energy with no energy input, sometimes known as a perpetual motion machine, cannot exist.

13 Energy Flow and Conservation of Energy Conservation of energy requires that the sum of the energy changes in the system and the surroundings must be zero. Energy universe = 0 = Energy system + Energy surroundings is the symbol that is used to mean change. Final amount initial amount

14 Internal Energy The internal energy ( E ) of a system is the sum of the kinetic and potential energies of all of the particles that compose the system The change in the internal energy of a system only depends on the amount of energy in the system at the beginning and end. A state function is a mathematical function whose result only depends on the initial and final conditions, not on the process used.

15 Energy Diagrams Energy diagram: compares the internal energy of the reactants and the products: The reactants become products, which have lower internal energy. Therefore, energy is given off by the reaction and E (that is, E products - E reactants ) is negative If we define the thermodynamic system as the reactants and products of the reaction

16 In this reaction E is positive and energy flows into the system and out of the surroundings.

17 Energy Flow in a Chemical Reaction In the reaction C(s) + O 2 (g) CO 2 (g), there will be a net release of energy into the surroundings. E reaction = E surroundings Internal Energy C(s), O 2 (g) CO 2 (g) energy released E rxn = Surroundings System C + O 2 CO 2

18 In the reaction CO 2 (g) C(s) + O 2 (g), there will be an absorption of energy from the surroundings into the reaction. Internal Energy C(s), O 2 (g) CO 2 (g) energy absorbed E rxn = + E reaction = E surroundings Surroundings System C + O 2 CO 2

19 Summary: Energy Flow Surroundings E System E + Surroundings E + System E

20 Energy Exchange Energy is exchanged between the system and surroundings through heat and work. q = heat (thermal) energy w = work energy q and w are NOT state functions; their value depends on the process.

21 TABLE 6.3 Sign Conventions for q, w, and ΔE

22 FIGURE 6.6 Energy, Work, and Heat

23 FIGURE 6.6 Energy, Work, and Heat

24 EXAMPLE 6.1 Internal Energy, Heat, and Work The firing of a potato cannon provides a good example of the heat and work associated with a chemical reaction. A potato is stuffed into a long cylinder that is capped on one end and open at the other. Some kind of fuel is introduced under the potato at the capped end usually through a small hole and ignited. The potato then shoots out of the cannon, sometimes flying hundreds of feet, and the cannon emits heat to the surroundings. If the burning of the fuel performs 855 J of work on the potato and produces 1422 J of heat, what is E for the burning of the fuel? (Note: A potato cannon can be dangerous and should not be constructed without proper training and experience.) Sol:

25 Heat Exchange Heat is the exchange of thermal energy between a system and surroundings. Heat exchange occurs when system and surroundings have a difference in temperature. Temperature is the measure of the thermal energy within a sample of matter. Heat flows from matter with high temperature to matter with low temperature until both objects reach the same temperature. Thermal equilibrium

26 Temperature Changes and Heat Capacity When a system absorbs heat ( q ), its temperature changes by ΔT : the heat absorbed by a system and its corresponding temperature change are directly proportional: q α ΔT. Units of C are J/ C or J/K

27 Heat Capacity the heat capacity ( C ) of a system as the quantity of heat required to change its temperature by 1 Heat capacity is an extensive property it depends on the amount of matter being heated specific heat capacity (Cs): the amount of heat required to raise the temperature of 1 gram of the substance by 1. molar heat capacity: the amount of heat required to raise the temperature of 1 mole of a substance by 1.

28 Heat Capacity The specific heat capacity of a substance can be used to quantify the relationship between the amount of heat added to a given amount of the substance and the corresponding temperature increase

29 EXAMPLE 6.2 Temperature Changes and Heat Capacity Suppose you find a penny (minted before 1982, when pennies were almost entirely copper) in the snow. How much heat is absorbed by the penny as it warms from the temperature of the snow, which is -8.0, to the temperature of your body, 37.0? Assume the penny is pure copper and has a mass of 3.10 g.

30 Heat Transfer and Final Temperature When two objects at different temperatures are placed in contact, heat flows from the material at the higher temperature to the material at the lower temperature. Heat flows until both materials reach the same final temperature. The amount of heat energy lost by the hot material equals the amount of heat gained by the cold material.

31 Thermal Energy Transfer A block of the metal to the water. metal at 55 C is added to water at 25 C. Thermal energy transfers heat from metal to water The exact temperature change depends on the following: The mass of the metal The mass of water Specific heat capacities of the metal and of water

32 EXAMPLE 6.3 Thermal Energy Transfer A 32.5 g cube of aluminum initially at 45.8 C is submerged into g of water at 15.4 C. What is the final temperature of both substances at thermal equilibrium? (Assume that the aluminum and the water are thermally isolated from everything else.)

33 Sol:

34 Work: Pressure Volume Work PV work is work caused by a volume change against an external pressure. When gases expand, V is positive, but the system is doing work on the surroundings, so w gas is negative. As long as the external pressure is kept constant, Work gas = External Pressure Change in Volume gas w = P V. To convert the units to joules use J = 1 atm L.

35 EXAMPLE 6.4 Pressure Volume Work To inflate a balloon you must do pressure volume work on the surroundings. If you inflate a balloon from a volume of L to 1.85 L against an external pressure of 1.00 atm, how much work is done (in joules)?

36 Exchanging Energy between System and Surroundings Exchange of heat energy q = mass specific heat Temperature Exchange of work w = Pressure Volume

37 Bomb Calorimeter It is used to measure E because it is a constant volume system. The heat capacity of the calorimeter is the amount of heat absorbed by the calorimeter for each degree rise in temperature and is called the calorimeter constant. (C cal, kj/ºc) the reaction occurs under conditions of constant volume

38 EXAMPLE 6.5 Measuring E rxn in a Bomb Calorimeter When g of sucrose (C 12 H 22 O 11 ) undergoes combustion in a bomb calorimeter, the temperature rises from C to Find E rxn for the combustion of sucrose in kj/mol sucrose. The heat capacity of the bomb calorimeter, determined in a separate experiment, is 4.90 kj/c. (You can ignore the heat capacity of the small sample of sucrose because it is negligible compared to the heat capacity of the calorimeter.)

39 Sol:

40 6.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure The enthalpy ( H ) of a system : as the sum of its internal energy and the product of its pressure and volume: H = E + PV H is a state function

41 The enthalpy change, Δ H, of a reaction is the heat evolved in a reaction at constant pressure. Δ H reaction = q reaction at constant pressure Usually DH and DE are similar in value; the difference is largest for reactions that produce or use large quantities of gas.

42 Endothermic and Exothermic Reactions Exothermic reaction : when ΔH is negative, heat is being released by the system. Endothermic reaction : When ΔH is positive, heat is being absorbed by the system.

43 EXAMPLE 6.6 Exothermic and Endothermic Processes Identify each process as endothermic or exothermic and indicate the sign of H. (a) sweat evaporating from skin (b) water freezing in a freezer (c) wood burning in a fire

44 Molecular View of Exothermic Reactions For an exothermic reaction, the surrounding s temperature rises due to a release of thermal energy by the reaction. This extra thermal energy comes from the conversion of some of the chemical potential energy in the reactants into kinetic energy in the form of heat. During the course of a reaction, existing bonds are broken and new bonds are made. The products of the reaction have less chemical potential energy than the reactants. The difference in energy is released as heat.

45 Molecular View of Endothermic Reactions In an endothermic reaction, the surrounding s temperature drops due to absorption of some of its thermal energy by the reaction. During the course of a reaction, existing bonds are broken and new bonds are made. The products of the reaction have more chemical potential energy than the reactants. To acquire this extra energy, some of the thermal energy of the surroundings is converted into chemical potential energy stored in the products.

46 Stoichiometry Involving H: Thermochemical Equations The enthalpy change for a chemical reaction, abbreviated H rxn, is also called the enthalpy of reaction or heat of reaction The enthalpy change in a chemical reaction is an extensive property. The more reactants you use, the larger the enthalpy change. when 1 mol of C 3 H 8 reacts with 5 mol of O 2 to form 3 mol of CO 2 and 4 mol of H 2 O, 2044 kj of heat is emitted

47 EXAMPLE 6.7 Stoichiometry Involving ΔH

48 Measuring H Calorimetry at Constant Pressure Reactions done in aqueous solution are at constant pressure. The calorimeter is often nested foam cups containing the solution. q reaction = - q solution = -(mass solution C s, solution T) H reaction = q constant pressure = q reaction To get H reaction per mol, divide by the number of moles.

49 EXAMPLE 6.8 Measuring H rxn in a Coffee-Cup Calorimeter

50 Sol:

51 Relationships Involving H rxn When reaction is multiplied by a factor, H rxn is multiplied by that factor. Because H rxn is extensive, C(s) + O 2 (g) CO 2 (g) H = kj 2 C(s) + 2 O 2 (g) 2 CO 2 (g) H = 2( kj) = kj. If a reaction is reversed, then the sign of H is changed. CO 2 (g) C(s) + O 2 (g) H = kj

52 Relationships Involving H rxn If a reaction can be expressed as a series of steps, then the H rxn for the overall reaction is the sum of the heats of reaction for each step. (Hess s Law)

53 EXAMPLE 6.9 Hess s Law

54 Sol:

55 6.9 Determining Enthalpies of Reaction from Standard Enthalpies of Formation The standard state is the state of a material at a defined set of conditions. Pure gas at exactly 1 atm pressure Pure solid or liquid in its most stable form at exactly 1 atm pressure and temperature of interest (Usually 25 C) Substance in a solution with concentration 1 M The standard enthalpy change, H, is the enthalpy change when all reactants and products are in their standard states.

56 Standard enthalpy of formation The standard enthalpy of formation, H f, is the enthalpy change for the reaction forming 1 mole of a pure compound from its constituent elements. The elements must be in their standard states. The H f for a pure element in its standard state = 0 kj/mol.

57 TABLE 6.5 Standard Enthalpies (or Heats) of Formation, H f o, at 298 K

58 Calculating the Standard Enthalpy Change for a Reaction The standard enthalpy of formation corresponds to the formation of a compound from its constituent elements in their standard states The decomposition of a compound into its constituent elements in their standard states:

59 Calculating Standard Enthalpy Change for a Reaction Any reaction can be written as the sum of formation reactions (or the reverse of formation reactions) for the reactants and products. The H for the reaction is then the sum of the H f for the component reactions. H reaction = n H f (products) n H f (reactants) means sum. n is the coefficient of the reaction.

60 EX:

61 Calculating H from Tabulated Values of H f

62 EXAMPLE 6.11 H o rxn and Standard Enthalpies of Formation Sol:

63 Energy Use and the Environment Most comes from the combustion of fossil fuels. Combustible materials that originate from ancient life. C(s) + O 2 (g) CO 2 (g) H rxn = kj CH 4 (g) +2 O 2 (g) CO 2 (g) + 2 H 2 O(g) H rxn = kj C 8 H 18 (g) O 2 (g) 8 CO 2 (g) + 9 H 2 O(g) H rxn = kj

64 Because of additives and impurities in the fossil fuel, incomplete combustion, and side reactions, harmful materials are added to the atmosphere when fossil fuels are burned for energy. Therefore, fossil fuel emissions contribute to air pollution, acid rain, and global warming.

65 Global Warming CO 2 is a greenhouse gas. It allows light from the sun to reach Earth, but does not allow the heat (infrared light) reflected off Earth to escape into outer space. CO 2 levels in the atmosphere have been steadily increasing. Current observations suggest that the average global air temperature has risen 0.6 C in the past 100 years.

66 Renewable Energy Our greatest unlimited supply of energy is the sun. New technologies are being developed to capture the energy of sunlight. Parabolic troughs, solar power towers, and dish engines concentrate the sun s light to generate electricity. Solar energy is used to decompose water into H 2 (g) and O 2 (g); the H 2 can then be used by fuel cells to generate electricity. H 2 (g) + ½ O 2 (g) H 2 O(l) H rxn = kj