大学化学课件-国外原版教材-5.ppt

Copyright 1999, PRENTICE HALL,Chapter 5,1,Thermochemistry,Chapter 5,David P. White University of North Carolina, Wilmington,Copyright 1999, PRENTICE HALL,Chapter 5,2,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:,Copyright 1999, PRENTICE HALL,Chapter 5,3,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 5,4,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 = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal,Copyright 1999, PRENTICE HALL,Chapter 5,5,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 5,6,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 5,7,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,First Law of Thermodynamics,Copyright 1999, PRENTICE HALL,Chapter 5,8,Relating DE to Heat and Work,First Law of Thermodynamics,Copyright 1999, PRENTICE HALL,Chapter 5,9,Relating DE to Heat and Work,First Law of Thermodynamics,Copyright 1999, PRENTICE HALL,Chapter 5,10,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,Copyright 1999, PRENTICE HALL,Chapter 5,11,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,Copyright 1999, PRENTICE HALL,Chapter 5,12,State Functions,First Law of Thermodynamics,Copyright 1999, PRENTICE HALL,Chapter 5,13,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,Enthalpy,Copyright 1999, PRENTICE HALL,Chapter 5,14,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,Enthalpies of Reaction,Copyright 1999, PRENTICE HALL,Chapter 5,15,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,Copyright 1999, PRENTICE HALL,Chapter 5,16,Heat Capacity and Specific Heat,Calorimetry,Copyright 1999, PRENTICE HALL,Chapter 5,17,Constant-Pressure Calorimetry Atmospheric pressure is constant! DH = qP qrxn = -qsoln = -(specific heat of solution) (grams of solution) DT.,Calorimetry,Copyright 1999, PRENTICE HALL,Chapter 5,18,Bomb Calorimetry (Constant-Volume Calorimetry) Reaction carried out under constant volume. Use a bomb calorimeter. Usually study combustion.,Calorimetry,Copyright 1999, PRENTICE HALL,Chapter 5,19,qrxn = -CcalorimeterT.,Bomb Calorimetry (Constant-Volume Calorimetry),Calorimetry,Copyright 1999, PRENTICE HALL,Chapter 5,20,Hesss 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: CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = -802 kJ 2H2O(g) 2H2O(l) H = -88 kJ CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = -890 kJ,Hesss Law,Copyright 1999, PRENTICE HALL,Chapter 5,21,In the above enthalpy diagram note that H1 = H2 + H3,Hesss Law,Copyright 1999, PRENTICE HALL,Chapter 5,22,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, Hof . Standard conditions (standard state): 1 atm and 25 oC (298 K). Standard enthalpy, Ho, 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 5,23,Enthalpies of Formation,Standard enthalpy of formation of the most stable form of an element is zero.,Copyright 1999, PRENTICE HALL,Chapter 5,24,Hrxn = H1 + H2 + H3,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.,Copyright 1999, PRENTICE HALL,Chapter 5,25,Enthalpies of Formation,Using Enthalpies of Formation to Calculate Enthalpies of Reaction For a reaction:,Copyright 1999, PRENTICE HALL,Chapter 5,26,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 fol