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Copyright 1999, PRENTICE HALLChapter 51 Thermochemistry David P. White University of North Carolina, Wilmington.

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Presentation on theme: "Copyright 1999, PRENTICE HALLChapter 51 Thermochemistry David P. White University of North Carolina, Wilmington."— Presentation transcript:

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2 Copyright 1999, PRENTICE HALLChapter 51 Thermochemistry David P. White University of North Carolina, Wilmington

3 Copyright 1999, PRENTICE HALLChapter 52 The Nature of Energy Kinetic and Potential Energy From Physics: Force is a push or pull on an object. Work is the product of force applied to an object over a distance: w = F  d Energy is the work done to move an object against a force. Kinetic energy is the energy of motion:

4 Copyright 1999, PRENTICE HALLChapter 53 The Nature of Energy Kinetic and Potential Energy Potential energy is the energy an object possesses by virtue of its position. Potential energy can be converted into kinetic energy. Example: a ball of clay dropping off a building.

5 Copyright 1999, PRENTICE HALLChapter 54 The Nature of Energy Energy Units SI Unit for energy is the joule, J: We sometimes use the calorie instead of the joule: 1 cal = J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal

6 Copyright 1999, PRENTICE HALLChapter 55 The Nature of Energy Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe.

7 Copyright 1999, PRENTICE HALLChapter 56 First Law of Thermodynamics Internal Energy Internal Energy: total energy of a system. Cannot measure absolute internal energy. Change in internal energy,  E = E final - E initial

8 Copyright 1999, PRENTICE HALLChapter 57 Relating  E to Heat and Work Energy cannot be created or destroyed. Energy of (system + surroundings) is constant. Any energy transferred from a system must be transferred to the surroundings (and vice versa). From the first law of thermodynamics: when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system:  E = q + w First Law of Thermodynamics

9 Copyright 1999, PRENTICE HALLChapter 58 Relating  E to Heat and Work First Law of Thermodynamics

10 Copyright 1999, PRENTICE HALLChapter 59 Relating  E to Heat and Work First Law of Thermodynamics

11 Copyright 1999, PRENTICE HALLChapter 510 Endothermic and Exothermic Processes Endothermic: absorbs heat from the surroundings. Exothermic: transfers heat to the surroundings. An endothermic reaction feels cold. An exothermic reaction feels hot. First Law of Thermodynamics

12 Copyright 1999, PRENTICE HALLChapter 511 State Functions State function: depends only on the initial and final states of system, not on how the internal energy is used. First Law of Thermodynamics

13 Copyright 1999, PRENTICE HALLChapter 512 State Functions First Law of Thermodynamics

14 Copyright 1999, PRENTICE HALLChapter 513 Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. Can only measure the change in enthalpy:  H = H final - H initial = q P Enthalpy

15 Copyright 1999, PRENTICE HALLChapter 514 For a reaction  H rxn = H(products) - H (reactants) Enthalpy is an extensive property (magnitude  H is directly proportional to amount): CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2CH 4 (g) + 4O 2 (g)  2CO 2 (g) + 4H 2 O(g)  H = kJ When we reverse a reaction, we change the sign of  H: CO 2 (g) + 2H 2 O(g)  CH 4 (g) + 2O 2 (g)  H = +802 kJ Change in enthalpy depends on state: H 2 O(g)  H 2 O(l)  H = -88 kJ Enthalpies of Reaction

16 Copyright 1999, PRENTICE HALLChapter 515 Heat Capacity and Specific Heat Calorimetry = measurement of heat flow. Calorimeter = apparatus that measures heat flow. Heat capacity = the amount of energy required to raise the temperature of an object (by one degree). Molar heat capacity = heat capacity of 1 mol of a substance. Specific heat = specific heat capacity = heat capacity of 1 g of a substance. q = (specific heat)  (grams of substance)  T. Be careful of the sign of q. Calorimetry

17 Copyright 1999, PRENTICE HALLChapter 516 Heat Capacity and Specific Heat Calorimetry

18 Copyright 1999, PRENTICE HALLChapter 517 Constant-Pressure Calorimetry Atmospheric pressure is constant!  H = q P q rxn = -q soln = -(specific heat of solution)  (grams of solution)   T. Calorimetry

19 Copyright 1999, PRENTICE HALLChapter 518 Bomb Calorimetry (Constant-Volume Calorimetry) Reaction carried out under constant volume. Use a bomb calorimeter. Usually study combustion. Calorimetry

20 Copyright 1999, PRENTICE HALLChapter 519 q rxn = -C calorimeter T. Bomb Calorimetry (Constant-Volume Calorimetry) Calorimetry

21 Copyright 1999, PRENTICE HALLChapter 520 Hess’s lawHess’s law: if a reaction is carried out in a number of steps,  H for the overall reaction is the sum of  H for each individual step. For example: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2H 2 O(g)  2H 2 O(l)  H = -88 kJ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ Hess’s Law

22 Copyright 1999, PRENTICE HALLChapter 521 In the above enthalpy diagram note that  H 1 =  H 2 +  H 3 Hess’s Law

23 Copyright 1999, PRENTICE HALLChapter 522 Enthalpies of Formation If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation,  H o f. Standard conditions (standard state): 1 atm and 25 o C (298 K). Standard enthalpy,  H o, is the enthalpy measured when everything is in its standard state. Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states. If there is more than one state for a substance under standard conditions, the more stable one is used.

24 Copyright 1999, PRENTICE HALLChapter 523 Enthalpies of Formation Standard enthalpy of formation of the most stable form of an element is zero.

25 Copyright 1999, PRENTICE HALLChapter 524  H rxn =  H 1 +  H 2 +  H 3 Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.

26 Copyright 1999, PRENTICE HALLChapter 525 Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction For a reaction:

27 Copyright 1999, PRENTICE HALLChapter 526 Foods and Fuels Foods Fuel value = energy released when 1 g of substance is burned. 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal. Energy in our bodies comes from carbohydrates and fats (mostly). Intestines: carbohydrates converted into glucose: C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O,  H = kJ Fats break down as follows: 2C 57 H 110 O O 2  114CO H 2 O,  H = -75,520 kJ Fats: contain more energy; are not water soluble, so are good for energy storage.

28 Copyright 1999, PRENTICE HALLChapter 527 Foods and Fuels Foods

29 Copyright 1999, PRENTICE HALLChapter 528 Foods and Fuels Fuels U.S.: 1.0 x 10 6 kJ of fuel per day. Most from petroleum and natural gas. Remainder from coal, nuclear, and hydroelectric. Fossil fuels are not renewable.

30 Copyright 1999, PRENTICE HALLChapter 529 Foods and Fuels Fuels Fuel value = energy released when 1 g of substance is burned. Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g.

31 Copyright 1999, PRENTICE HALLChapter 530 End of Chapter 5 Thermochemistry


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