Chapter 5 - Chapter 5 Thermochemistry General Chemistry I,...

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Unformatted text preview: Chapter 5 Thermochemistry General Chemistry I, 2009 Professor Maggie Ciszkowska Thermochemistry The Nature of Energy Chemical reaction: A+BC+D Usually (not always), there is one more product of this reaction - ENERGY Reaction can be exothermic heat is produced by the system (rxn) Reaction can be endothermic heat is consumed by the system (rxn) (taken from the surroundings) General Chemistry I, 2009 Professor Maggie Ciszkowska Thermochemistry The Nature of Energy Kinetic Energy and Potential Energy Kinetic energy is the energy of motion: 1 2 Ek = mv 2 Potential energy is the energy an object possesses by virtue of its position; PE = mgh m = mass, g = gravitational acceleration, h = height. Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill. General Chemistry I, 2009 Professor Maggie Ciszkowska The Nature of Energy Units of Energy SI Unit for energy is the joule, J: 1 2 1 2 Ek = mv = (2 kg )(1 m/s ) 2 2 2 2 = 1 (kg m /s = 1 J We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) Note : A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal General Chemistry I, 2009 Professor Maggie Ciszkowska The Nature of Energy Work Work = forcedistance work = F d = (1N )(1m ) = 1 kg(m /s ) 2 2 General Chemistry I, 2009 Professor Maggie Ciszkowska The Nature of Energy Energy and Heat The energy (kinetic and potential) can be converted into heat. When you warm up an object its energy increases. At the molecular level: molecules are moving faster at higher energy; they have higher kinetic energy; their potential energy is higher. Energy at molecular level is called "internal energy"; energy that we are interested in chemistry. General Chemistry I, 2009 Professor Maggie Ciszkowska The Nature of Energy Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe. The system interacts with surroundings. We are interested in changes of internal energy of the system. Closed system a system when we do not add or remove anything. General Chemistry I, 2009 Professor Maggie Ciszkowska The Nature of Energy Transferring Energy: Work and Heat 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 1 J = 1 Nm = kg(m2/s2) Energy is the work done to move an object against a force. Heat is the transfer of energy between two objects. Energy is the capacity to do work or transfer heat. General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics Internal Energy Internal Energy: total energy of a system. Cannot measure absolute internal energy. Change in internal energy, E = Efinal - Einitial General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics Relating E to Heat and Work Energy cannot be created or destroyed (conservation of energy). 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 General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics Conservation of Energy E ( system) + E ( surroundings ) = 0 General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics General Chemistry I, 2009 Professor Maggie Ciszkowska General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics Calculate the change of internal energy of a system, if 0.010 kJ of heat is added to this system and the system does 20 J of work. General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics Calculate the change of internal energy of a system, if 0.010 kJ of heat is added to this system and the system does 20 J of work. E = q + w q = 10 J w = -20 J E = 10 J + (-20 J ) = -10 J General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics Exothermic and Endothermic Processes Endothermic: system absorbs heat from the surroundings. Exothermic: system transfers heat to the surroundings. An endothermic reaction feels cold. An exothermic reaction feels hot. General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics State Functions State function property of the system: depends only on the initial and final states of system, not on how the internal energy is used. General Chemistry I, 2009 Professor Maggie Ciszkowska The First Law of Thermodynamics State Functions State function property of the system: depends only on the initial and final states of system, not on how the internal energy is used. P, V, T, composition, internal energy General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Chemical reactions can absorb or release heat. However, they also have the ability to do work. For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing work. Zn(s) + 2H+(aq) Zn2+(aq) + H2(g) The work performed by the above reaction is called pressure-volume work. When the pressure is constant, P=const., w = - P V General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. H = E + PV Enthalpy is a state function. If the process occurs at constant pressure, H = (E + PV ) H = E + PV General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Since we know that w = - PV We can write H = E + PV H = q P + w When H is positive, the system gains heat from the surroundings endothermic process. When H is negative, the surroundings gain heat from the system exothermic process. General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Since we know that when P=const.: w = - PV and that H = E + PV = qP + w Therefore, under constant pressure: H = E + PV = E - w H = (qP + w) - w = qP General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Additionally, for most of chemical reactions, the magnitude of heat, qp, is much larger than work, w: E = (qP + w) = qP Therefore, the change of enthalpy is very close to the change of internal energy: H = E General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Indicate the sign of the enthalpy change, H, in each of the following processes, P = const.: a. an ice cube melts b. 1 g of C4H10 is combusted in O2 to produce CO2 and H2O. General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpy Indicate the sign of the enthalpy change, H, in each of the following processes, P=const.: a. an ice cube melts The ice cube (system) absorbs heat from the surroundings. H>0, endothermic process b. 1 g of C4H10 is combusted in O2 to produce CO2 and H2O. This reaction produces heat, the energy in a form of heat flows from the system to the surrounding. H<0, exothermic process General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions For a reaction: H = H final - H initial = H products - H reactants Enthalpy is an extensive property (magnitude H is directly proportional to the amount of reacting compounds): CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = -890 kJ 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) H = -1780 kJ General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions When we reverse a reaction, we change the sign of H: CH4(g) + 2O2(g) CO2(g) + 2H2O(g) CO2(g) + 2H2O(g) CH4(g) + 2O2(g) H = -890 kJ H = +890 kJ Change in enthalpy depends on a state: H2O(g) H2O(l) H = -88 kJ General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions What is called H is really H - standard enthalpy (at 25 C (298 K) and 1 atm) General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions What is the change of internal energy in an insulated system when 1 mol of CH4(g) is burned with 2 mol of O2(g) at 25 C (this is an exothermic reaction). E > 0, E < 0, or E = 0 ? General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions What is the change of internal energy in an insulated system when 1 mol of CH4(g) is burned with 2 mol of O2(g) at 25 C (this is an exothermic reaction). E > 0, E < 0, or E = 0 ? Insulated system does not exchange energy with the surroundings. E = 0 General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions Use H = -890 kJ to find the heat produced when 4.80 g of CH4 gas is burned in oxygen at 25 C and 1 atm. General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions Use H = -890 kJ to find the heat produced when 4.80 g of CH4 gas is burned in oxygen at 25 C and 1 atm. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = -890 kJ 1 mol 4.80 g = 0.300mol 16.0 g - 890 kJ 0.300 mol = -267 kJ 1 mol General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions If H = -890 kJ for the following chemical reaction: 1/2 CH4(g) + O2(g) 1/2 CO2(g) + H2O(g) calculate H for this process: 1/2 CH4(g) + O2(g) 1/2 CO2(g) + H2O(g) H = ? kJ General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Chemical Reactions If H = -890 kJ for the following chemical reaction: 1/2 CH4(g) + O2(g) 1/2 CO2(g) + H2O(g) calculate H for this process: 1/2 CH4(g) + O2(g) 1/2 CO2(g) + H2O(g) H = -445 kJ General Chemistry I, 2009 Professor Maggie Ciszkowska Calorimetry Heat Capacity and Specific Heat Calorimetry = measurement of heat flow. Calorimeter = apparatus that measures heat flow. Heat is transferred from one system to another system. This heat is directly proportional to the mass of a heated system, and to the change of temperature of this system. Specific heat = specific heat capacity = heat capacity of 1 g of a substance. q = (specific heat ) (mass) T q = (s ) (m ) T General Chemistry I, 2009 Professor Maggie Ciszkowska Calorimetry Specific Heat Specific heat s = q / (mT) Units: J/(gK) = Jg-1K-1 For example, for H2O: s = 4.18 Jg-1K-1 General Chemistry I, 2009 Professor Maggie Ciszkowska Calorimetry Constant Pressure Calorimetry Atmospheric pressure is constant. q = (s ) (m ) T H = qP qrxn = -qsoln = -(specific heat of solution ) (grams of solution ) T General Chemistry I, 2009 Professor Maggie Ciszkowska Calorimetry Constant Pressure Calorimetry General Chemistry I, 2009 Professor Maggie Ciszkowska Calorimetry When 50 mL of 0.1 M HCl(aq) is mixed with 50 mL of 0.10 M NaOH(aq) in an insulated container, the temperature rises from 21.0 C to 27.5 C. Find the H for this reaction. Is this reaction exothermic or endothermic? Assume: s = 4.18 Jg-1K-1 d = 1.0 gmL-1 General Chemistry I, 2009 Professor Maggie Ciszkowska When 50.0 mL of 0.1 M HCl(aq) is mixed with 50.0 mL of 0.10 M NaOH(aq) in an insulated container, the temperature rises from 21.0 C to 27.5 C; P = const. Find the H for this reaction. Is this reaction exothermic or endothermic? Assume: s = 4.18 Jg-1K-1 and d = 1.00 gmL-1 HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) m (in grams) = (50.0 mL + 50.0 mL)1.00 g/mL = 100 g T = 27.5 C - 21.0 C = 6.5 C = 6.5 K For diluted solutions: H = -smT H = -(4.18 Jg-1K-1 100 g 6.5 K) = -2700 J = -2.7 kJ Since the process occurs at constant pressure, we have: H = qp = -2.7 kJ exothermic process General Chemistry I, 2009 Professor Maggie Ciszkowska Hess's Law Hess'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. H depends only on the amount of substance that undergo changes (extensive property), and the final and initial state of the system, products and reactants (state function). 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 General Chemistry I, 2009 Professor Maggie Ciszkowska Hess's Law Note that: H1 = H2 + H3 General Chemistry I, 2009 Professor Maggie Ciszkowska Hess's Law Knowing the following enthalpies of reactions: H = -393.5 kJ/mol C C(s) + O2(g) CO2(g) CO(g) + 1/2 O2(g) CO2(g) H = -283.0 kJ/mol CO calculate H for: C (s) + O2(g) CO (g) General Chemistry I, 2009 Professor Maggie Ciszkowska Hess's Law Knowing the following enthalpies of reactions: H = -393.5 kJ/mol C C(s) + O2(g) CO2(g) CO(g) + 1/2 O2(g) CO2(g) H = -283.0 kJ/mol CO calculate H for: C (s) + O2(g) CO (g) C(s) + O2(g) CO2(g) CO2(g) CO(g) + O2(g) C(s) + O2(g) CO(g) H = -393.5 kJ + 283.0 kJ = -110.5 kJ General Chemistry I, 2009 Professor Maggie Ciszkowska H = -393.5 kJ H = +283.0 kJ Hess's Law Calculate H for: 2 C (s) + H2(g) C2H2 (g) given the following chemical eqns and H values: C2H2(g) + 5/2 O2(g) 2 CO2(g) + H2O(l) C(s) + O2(g) CO2(g) H2(g) + O2(g) H2O(l) H = -1299.6 kJ H = -393.5 kJ H = -285.8 kJ General Chemistry I, 2009 Professor Maggie Ciszkowska Calculate H for: 2 C (s) + H2(g) C2H2 (g) given the following chemical eqns and H values: C2H2(g) + 5/2 O2(g) 2 CO2(g) + H2O(l) H = -1299.6 kJ C(s) + O2(g) CO2(g) H = -393.5 kJ H2(g) + O2(g) H2O(l) H = -285.8 kJ 2 CO2(g) + H2O(l) C2H2(g) + 5/2 O2(g) 2 C(s) + 2 O2(g) 2 CO2(g) H2(g) + O2(g) H2O(l) 2 C(s) + H2(g) C2H2(g) H = +1299.6 kJ H = 2(-393.5 kJ) H = -285.8 kJ H = +1299.6 kJ + 2(-393.5 kJ) + (-285.8 kJ) = 226.8 kJ endothermic process General Chemistry I, 2009 Professor Maggie Ciszkowska 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 . (subscript f formation) Standard conditions (std. state): 1 atm and 25oC (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. General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation If there is more than one state for a substance under standard conditions, the more stable one is used. Standard enthalpy of formation of the most stable form of an element is zero. General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation 2 C(graphite) + 3 H2(g) + O2(g) C2H5OH(l) Hf= -277.7 kJ H2(g) + O2(g) H2O(l) Hf = -285.8 kJ O2(g) O2(g) Hf = 0 Hf for elements = 0 General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Write the formation reaction for NH3, characterized by Hf for NH3(g). General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Write the formation reaction for NH3, characterized by Hf for NH3(g). 1/2 N2(g) + 3/2 H2(g) NH3(g) General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Using Enthalpies of Formation of Calculate Enthalpies of Reaction For a reaction H rxn = nH f (products ) - mH f (reactants ) m and n numbers of moles General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Find the enthalpy of the following reaction, Hrxn (standard, 298 K): aA+bBcC+dD knowing only standard enthalpies of formation: Hf (A) Hf (B) Hf (C) Hf (D) General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Find the enthalpy of the following reaction, Hrxn (standard, 298 K): aA+bBcC+dD knowing only standard enthalpies of formation: Hf (A) Hf (B) Hf (C) Hf (D) Hrxn = [(cHf(C) + dHf(D)] - [aHf(A) + bHf(B)] General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Find the enthalpy of the following reaction, Hrxn (standard, 298 K): 2 C2H6(g) + 7 O2(g) 4 CO2(g) + 6 H2O(l) knowing only standard enthalpies of formation: Hf (for CO2(g)) = - 393.5 kJ/mol Hf (for H2O(l)) = -286 kJ/mol Hf (for C2H6(g)) = -85 kJ/mol Hf (for O2(g)) = 0 kJ/mol General Chemistry I, 2009 Professor Maggie Ciszkowska Enthalpies of Formation Find the enthalpy of the following reaction, Hrxn (standard, 298 K): 2 C2H6(g) + 7 O2(g) 4 CO2(g) + 6 H2O(l) knowing only standard enthalpies of formation: Hrxn = [(4Hf(CO2(g)) + 6Hf(H2O(l))] - -[2Hf(C2H6(g)) + 7Hf(O2(g))] Hrxn = [4(-393.5 kJ/mol) + 6(-286 kJ/mol)] - - [2(-85 kJ/mol) + 7(0 kJ/mol)] = -3120 kJ/mol General Chemistry I, 2009 Professor Maggie Ciszkowska 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). Carbohydrates converted into glucose and then burned: C6H12O6 + 6O2 6CO2 + 6H2O, H = -2,816 kJ H = -75,520 kJ Fats break down as follows: 2C57H110O6 + 163O2 114CO2 + 110H2O, General Chemistry I, 2009 Professor Maggie Ciszkowska Fuels In 2000 the United States consumed 1.03 1017 kJ of fuel. Most from petroleum and natural gas. Remainder from coal, nuclear, and hydroelectric. Fossil fuels are not renewable. General Chemistry I, 2009 Professor Maggie Ciszkowska Fuels General Chemistry I, 2009 Professor Maggie Ciszkowska 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. General Chemistry I, 2009 Professor Maggie Ciszkowska ...
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This note was uploaded on 05/04/2009 for the course CHEM 0207 taught by Professor Cizkowska during the Spring '09 term at CUNY Brooklyn.

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