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Thermochemistry David P. White University of North Carolina, Wilmington Chapter 5 Copyright 1999, PRENTICE HALL Chapter 5 1 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=Fd • Energy is the work done to move an object against a force. • Kinetic energy is the energy of motion: 2 1 E k mv 2 Copyright 1999, PRENTICE HALL Chapter 5 2 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. Copyright 1999, PRENTICE HALL Chapter 5 3 The Nature of Energy Energy Units SI Unit for energy is the joule, J: E k 1 mv 2 1 2 kg 1 m/s 2 2 2 1 kg m 2 / s 2 1J We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal Copyright 1999, PRENTICE HALL Chapter 5 4 The Nature of Energy Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe. Copyright 1999, PRENTICE HALL Chapter 5 5 First Law of Thermodynamics Internal Energy • Internal Energy: total energy of a system. • Cannot measure absolute internal energy. • Change in internal energy, DE = Efinal Einitial Copyright 1999, PRENTICE HALL Chapter 5 6 First Law of Thermodynamics Relating DE 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: DE = q + w Copyright 1999, PRENTICE HALL Chapter 5 7 First Law of Thermodynamics Relating DE to Heat and Work Copyright 1999, PRENTICE HALL Chapter 5 8 First Law of Thermodynamics Relating DE to Heat and Work Copyright 1999, PRENTICE HALL Chapter 5 9 First Law of Thermodynamics 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. Copyright 1999, PRENTICE HALL Chapter 5 10 First Law of Thermodynamics State Functions State function: depends only on the initial and final states of system, not on how the internal energy is used. Copyright 1999, PRENTICE HALL Chapter 5 11 First Law of Thermodynamics State Functions Copyright 1999, PRENTICE HALL Chapter 5 12 Enthalpy Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. Can only measure the change in enthalpy: DH = Hfinal - Hinitial = qP Copyright 1999, PRENTICE HALL Chapter 5 13 Enthalpies of Reaction For a reaction DHrxn = H(products) - H (reactants) Enthalpy is an extensive property (magnitude DH is directly proportional to amount): CH4(g) + 2O2(g) CO2(g) + 2H2O(g) DH = -802 kJ 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) DH = -1604 kJ When we reverse a reaction, we change the sign of DH: CO2(g) + 2H2O(g) CH4(g) + 2O2(g) DH = +802 kJ Change in enthalpy depends on state: H2O(g) H2O(l) DH = -88 kJ Copyright 1999, PRENTICE HALL Chapter 5 14 Calorimetry 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. Copyright 1999, PRENTICE HALL Chapter 5 15 Calorimetry Heat Capacity and Specific Heat Copyright 1999, PRENTICE HALL Chapter 5 16 Calorimetry Constant-Pressure Calorimetry Atmospheric pressure is constant! DH = qP qrxn = -qsoln = -(specific heat of solution) (grams of solution) DT. Copyright 1999, PRENTICE HALL Chapter 5 17 Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry) Reaction carried out under constant volume. Use a bomb calorimeter. Usually study combustion. Copyright 1999, PRENTICE HALL Chapter 5 18 Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry) qrxn = -CcalorimeterT. Copyright 1999, PRENTICE HALL Chapter 5 19 Hess’s Law • Hess’s law: if a reaction is carried out in a number of steps, DH for the overall reaction is the sum of DH for each individual step. • For example: CH4(g) + 2O2(g) CO2(g) + 2H2O(g) DH = -802 kJ 2H2O(g) 2H2O(l) DH = -88 kJ CH4(g) + 2O2(g) CO2(g) + 2H2O(l) DH = -890 kJ Copyright 1999, PRENTICE HALL Chapter 5 20 Hess’s Law In the above enthalpy diagram note that DH1 = DH2 + DH3 Copyright 1999, PRENTICE HALL Chapter 5 21 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, DHof . • Standard conditions (standard state): 1 atm and 25 oC (298 K). • Standard enthalpy, DHo, 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. Copyright 1999, PRENTICE HALL Chapter 5 22 Enthalpies of Formation • Standard enthalpy of formation of the most stable form of an element is zero. Copyright 1999, PRENTICE HALL Chapter 5 23 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. DHrxn = DH1 + DH2 + DH3 Copyright 1999, PRENTICE HALL Chapter 5 24 Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction For a reaction: DH rxn nDH f products mDH f reactants Copyright 1999, PRENTICE HALL Chapter 5 25 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: C6H12O6 + 6O2 6CO2 + 6H2O, DH = -2816 kJ •Fats break down as follows: 2C57H110O6 + 163O2 114CO2 + 110H2O, DH = -75,520 kJ •Fats: contain more energy; are not water soluble, so are good for energy storage. Copyright 1999, PRENTICE HALL Chapter 5 26 Foods and Fuels Foods Copyright 1999, PRENTICE HALL Chapter 5 27 Foods and Fuels Fuels U.S.: 1.0 x 106 kJ of fuel per day. Most from petroleum and natural gas. Remainder from coal, nuclear, and hydroelectric. Fossil fuels are not renewable. Copyright 1999, PRENTICE HALL Chapter 5 28 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. Copyright 1999, PRENTICE HALL Chapter 5 29 Thermochemistry End of Chapter 5 Copyright 1999, PRENTICE HALL Chapter 5 30