1 THERMOCHEMISTRY Engr. Edgie L. Estopace School of Chemical Engineering and Chemistry Mapua Institute of Technology Outline 1. Principles of heat flow 2. Measurement of heat flow; calorimetry 3. Enthalpy 4. Thermochemical equations 5. Enthalpies of formation 6. Bond enthalpy 7. The first law of thermodynamics Principles of Heat Flow Definitions The system: that part of the universe on which attention is focused The surroundings: the rest of the universe Practically speaking, it is possible to consider only the surroundings that directly contact the system Figure 6.1 Systems and Surroundings Chemical Reactions When we study a chemical reaction, we consider the system to be the reactants and products The surroundings are the vessel (beaker, test tube, flask) in which the reaction takes place plus the air or other material in thermal contact with the reaction system Forms of Energy Two broad categories of energy: potential energy and kinetic energy. Potential energy - associated with the relative position of an object. Kinetic energy - associated with motion. Kinetic energy = 12 mv2Forms of Energy Internal energy - the combined kinetic and potential energies of atoms and molecules that make up an object or system. Chemical energy - energy released or absorbed during a chemical reaction. Other forms of energy include radiant, mechanical, thermal, electrical, and nuclear. Thermochemistry - the study of the energetic consequences of chemistry Heat and Work Heat is the flow of energy between two objects because of a difference in temperature. Heat flows from the warmer object to the cooler object. Work is the transfer of energy accomplished by a force moving a mass some distance against resistance. Pressure-volume work (PV-work) is the most common work type in chemistry. Releasing an inflated balloon before it is tied off illustrates an example of PV-work. State Properties The state of a system is specified by enumerating: Composition Temperature Pressure State properties depend only on the state of the system, not on the path the system took to reach the state Mathematically for a state property X: X is the change in X X = Xfinal Xinitial Direction and Sign of Heat Flow Heat is given the symbol, q q is positive when heat flows into the system from the surroundings q is negative when heat flows from the system into the surroundings Endothermic processes have positive q H2O (s) H2O () q > 0 Exothermic processes have negative q CH4 (g) + 2O2 (g) CO2 (g) + H2O () q < 0 Exothermic and Endothermic Processes Magnitude of Heat Flow In any process, we are interested in both the direction of heat flow and in its magnitude q is expressed in joules (or kilojoules) James Joule (1818-1889); calorimetry Alternate unit: calorie 1 calorie = 4.184 J 1 kilocalorie = 4.184 kJ Nutritional calories are kcal The Calorimetry Equation q = C x t t = tfinal tinitial C (uppercase) is the heat capacity of the system: it is the quantity of heat needed to raise the temperature of the system by 1 C q = m x c x t c (lowercase) is the specific heat: the quantity of heat needed to raise the temperature of one gram of a substance by 1 C c depends on the identity and phase of the substance Heat Capacity and Specific Heat The amount of heat energy absorbed can be quantified. q = mcATq = nCpATHeat Capacity and Specific Heat Specific heat and molar heat capacities for some common substances. Example Problem 2 Heating a 24.0 g aluminum can raises its temperature by 15.0C. Find the value of q for the can. Example Problem 3 The molar heat capacity of liquid water is 75.3 J/mol K. If 37.5 g of water is cooled from 42.0 to 7.0C, what is q for the water? Specific Heat The specific heat of a substance, like the density or melting point, is an intensive property that can be used to identify a substance or determine its purity Water Water has an unusually large specific heat A large quantity of heat is required to raise the temperature of water Climate is moderated by the specific heat of water Only two states in the US have never recorded temperatures over 100 F one is Alaska (cold North) and the other is Hawaii (moderated by water) Table 6.1 Example 6.1 Measurement of Heat Flow: Calorimetry A calorimeter is a device used to measure the heat flow of a reaction The walls of the calorimeter are insulated to block heat flow between the reaction and the surroundings The heat flow for the system is equal in magnitude and opposite in sign from the heat flow of the calorimeter qreaction = - qcalorimeter qreaction = - Ccal t Calorimetry Heat flow is measured using a calorimeter. A calorimeter measures the heat evolved or absorbed by the system of interest by measuring the temperature change in the surroundings. qsystem = qsurroundingsqgained = qlostExample Problem 9.4 A glass contains 250.0 g of warm water at 78.0C. A piece of gold at 2.30C is placed in the water. The final temperature reached by this system is 76.9C. What was the mass of gold? The specific heat of water is 4.184 J/gC, and that of gold is 0.129 J/gC. Calorimetry There are two steps in a calorimetric measurement. Calibration - the calorimeter constant, Ccalorimeter, is determined by dividing the known amount of heat released in the calorimeter by the temperature change of the calorimeter. Actual Measurement - heat released or absorbed in a reaction of known quantity of material is measured. q = Ccalorimeter ATCalorimetry Actual Measurement - temperature change for the calorimeter and the calorimeter constant are used to determine the amount of heat released by a reaction. qcalorimeter = Ccalorimeter ATcalorimeterqreaction = qcalorimeterCalorimetry Diagram of a bomb calorimeter and standard choice for system and surroundings in a bomb calorimetry experiment. Example Problem 9.5 In the calibration of a calorimeter, an electrical resistance heater supplies 100.0 J of heat and a temperature increase of 0.850C is observed. Then, 0.245 g of a particular fuel is burned in this same calorimeter and the temperature increases by 5.23C. Calculate the energy density of this fuel, which is the amount of energy liberated per gram of fuel burned. Figure 8.2 Coffee-cup Calorimeter For a reaction performed in a coffee-cup calorimeter tC gJm q water reaction A = 18 . 4Example 6.2 Example 6.2, (Contd) Figure 6.3 Bomb Calorimeter The bomb calorimeter is more versatile than the coffee-cup calorimeter Reactions involving high temperature Reactions involving gases The bomb is a heavy metal vessel that is usually surrounded by water qreaction = -qcalorimeter qreaction = -Ccalt Ccal is a function of the calorimeter and can be measured experimentally Example 8.3 Enthalpy The heat flow at constant pressure is equal to the difference in enthalpy (heat content) between products and reactants The symbol for enthalpy is H We measure changes in enthalpy using a calorimeter and a reaction run at constant pressure: H = Hproducts Hreactants The sign of the enthalpy change is the same as for heat flow: H > 0 for endothermic reactions H < 0 for exothermic reactions Enthalpy is a state variable Enthalpy The conditions under which heat flow, q, occurs will have an impact on the measurement that is made. Combustion of octane releases 5.45 x 103 kJ under constant volume conditions, represented as qv. Combustion of octane releases 5.48 x 103 kJ under constant pressure conditions, represented as qp. Defining Enthalpy The internal energy change for a reaction equals the sum of the heat flow and the work. During an expansion, w = PAV. Under constant volume conditions, AV = 0, and AE = qv. AE = q + wAE = q PAVAE = qvDefining Enthalpy Enthalpy is the heat flow under conditions of constant pressure. The enthalpy change can be expressed as H = E + PV AH = AE + A(PV)AH = qpAH = (q PAV) + A(PV)AH = q PAV + PAVDefining Enthalpy When a system releases heat, the process is said to be exothermic. The value of AH is less than zero; the sign on AH is negative When a system absorbs heat, the process is said to be endothermic. The value of AH is greater than zero; the sign on AH is positive. AH of Phase Changes Phase changes occur under constant pressure conditions. The heat flow during a phase change is an enthalpy change. During a phase change, temperature does not change with heat flow due to formation or breaking of intermolecular attractive forces. H of Phase Changes The heat required to convert a liquid to a gas is the heat of vaporization, AHvap. AHvap is endothermic with a positive value. The heat released to convert a gas to a liquid is the heat of condensation, AHcond. AHcond is exothermic with a negative value. AHcond = AHvap The values of enthalpy changes in opposite directions have equal numeric values and differ only in their signs. The magnitude of enthalpy change depends on the substance involved. H of Phase Changes Standard molar enthalpies and temperatures for phase changes of water. H of Phase Changes The value of AH for a phase change is compound specific and has units of kJ/mol. The heat flow can be calculated using the number of moles of substance, n, and the value of the enthalpy change. AH = n AHphase changeH of Phase Changes Heat curve for the heating of 500-g of ice at -50oC to 200oC. Example Problem Calculate the enthalpy change when 240. g of ice melts. Bonds and Energy The enthalpy change for a reaction can be estimated using bond energies. During a chemical reaction, reactant bonds are broken and product bonds are made. Breaking bonds requires energy. Making bonds releases energy. If the amount of energy released making product bonds is greater than the amount of energy required to break reactant bonds, the reaction is exothermic. If the energy released is less than the energy required, the reaction is endothermic. Bonds and Energy The combustion of methane breaks 4 C-H bonds and 2 O=O bonds. 2 C=O bonds and 4 O-H bonds are made. CH4 + 2O2 CO2 + 2H2OBonds and Energy The accuracy of enthalpy changes calculated from tabulated bond energies is not very good. The bond energies used are averages. Bond energy method used to estimate enthalpy changes for reactions involving compounds with no available thermochemical data. A thermochemical equation summarizes the overall energetics for a chemical reaction. The sign on the AH indicates whether the reaction is endothermic or exothermic CH4(g) + 2O2(g) CO2(g) + 2H2O(l )AH = 890.4 kJBonds and Energy The combustion of methane is an exothermic reaction and releases 890.4 kJ of heat energy when 1 mole of methane reacts with 2 moles of oxygen. For thermochemical equations, if the stoichiometric coefficients are multiplied by some factor, the heat of reaction must also be multiplied by the same factor. CH4(g) + 2O2(g) CO2(g) + 2H2O(l )AH = 890.4 kJ2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(l )AH = 1780.8 kJHeats of Reaction for Some Specific Reactions Some classes of chemical reactions are given their own labels for heats of reactions. Heat of combustion, AHcomb Heat of neutralization, AHneut Heat of formation, AHf, is the heat of reaction for formation of substances. Fractional coefficients are allowed for formation reactions because only one mole of product can be formed. C(s) + 12 O2(g) CO(g)Heats of Reaction for Some Specific Reactions A formation reaction is the chemical reaction by which one mole of a compound is formed from its elements in their standard states. The standard state is the most stable form of an element at room temperature, 25oC, and pressure, 1 atm, indicated with a superscript o. AHfo = 0 for an element in its standard state. Hesss Law and Heats of Reaction Direct calorimetric determinations of some reactions may be too difficult or dangerous to perform. An indirect method is needed to obtain heats of reaction. Hesss law: the enthalpy change for any process is independent of the particular way the process is carried out. Enthalpy is a state function. A state function is a variable whose value depends only on the state of the system and not its history. Hesss Law Conceptual diagram representing Hesss law. Enthalpy is a state function, so any convenient path can be used to calculate the enthalpy change. Hesss Law Enthalpy diagram for the combustion of methane. The CH4 is converted to CO, then the CO is converted to CO2. The AH for each step is used to calculate the AH for the overall reaction. The AH will be the same for both paths. Example Problem One origin of SO3 is the combustion of sulfur, which is present in small quantities in coal, according to the following equation. Given the thermochemical information below, determine the heat of reaction for this reaction. S(s) +32O2(g) SO3(g)S(s) + O2(g) SO2(g) AHo= 296.8 kJ2SO2(g) + O2(g) 2SO3(g) AHo= 197.0 kJFormation Reactions and Hesss Law Conceptual diagram showing how to use tabulated enthalpies of formation to calculate the enthalpy change for a chemical reaction. Formation Reactions and Hesss Law The enthalpy change for a reaction can be calculated using Hesss law and heats of formation. AHo= viAHfoi(products)i vjAHfoj(reactants)jExample Problem 9.8 Use tabulated data to find the heat of combustion of one mole of propane, C3H8, to form gaseous carbon dioxide and liquid water. Example Problem 9.9 Ethanol, C2H5OH, is used to introduce oxygen into some blends of gasoline. It has a heat of combustion of 1366.8 kJ/mol. What is the heat of formation of ethanol? Energy and Stoichiometry A thermochemical equation allows for the stoichiometric treatment of energy. For an exothermic reaction, energy is treated as a product. For an endothermic reaction, energy is treated as a reactant. The thermochemical equation is used to convert between the number of moles of a reactant or product and the amount of energy released or absorbed. The stated value of AH for a thermochemical equation corresponds to the reaction taken place exactly as written, with the indicated numbers of moles of each substance reacting. Energy and Stoichiometry Flow chart detailing the sequence of steps needed to calculate the amount of energy released or absorbed when a chemical reaction is carried out using a given amount of material. Thermochemical Equations A thermochemical equation is a chemical equation with the H for the reaction included Example NH4NO3 (s) NH4+ (aq) + NO3- (aq) Experiment gives qreaction = 351 J for one gram of ammonium nitrate For one mole, this is The thermochemical equation is NH4NO3 (s) NH4+ (aq) + NO3- (aq) H = +28.1 kJ kJ J XmolggJ1 . 28 10 81 . 2105 . 8000 . 13514= = Conventions for Thermochemical Equations 1. The sign of H indicates whether the reaction is endothermic or exothermic 2. The coefficients of the thermochemical equation represent the number of moles of reactant and product 3. The phases of all reactant and product species must be stated 4. The value of H applies when products and reactants are at the same temperature, usually 25 C Rules of Thermochemistry 1. The magnitude of H is directly proportional to the amount of reactant or product 2. H for the reaction is equal in magnitude but opposite in sign for H for the reverse of the reaction 3. The value of H is the same whether the reaction occurs in one step or as a series of steps This rule is a direct consequence of the fact that H is a state variable This rule is a statement of Hesss Law Example 6.4 Example 6.4, (Contd) Enthalpy of Phase Changes Phase changes involve enthalpy There is no change in temperature during a phase change Endothermic: melting or vaporization Exothermic: freezing or condensation Pure substances have a value of H that corresponds to melting (reverse, fusion) or vaporization (reverse, condensation) Example 6.5 Example 6.6 Example 6.6, (Contd) Recap of the Rules of Thermochemistry H is directly proportional to the amount of reactant or product If a reaction is divided by 2, so is H If a reaction is multiplied by 6, so is H H changes sign when the reaction is reversed H has the same value regardless of the number of steps Enthalpies of Formation The standard molar enthalpy of formation, , is equal to the enthalpy change For one mole of a compound At constant pressure of 1 atm At a fixed temperature of 25 C From elements in their stable states at that temperature and pressure Enthalpies of formation are tabulated in Table 8.3 and in Appendix 1 in the back of the textbook fH ATable 6.3 Table 6.3, (Contd) Table 6.3, (Contd) Enthalpies of Formation of Elements and of H+ (aq) The enthalpy of formation of an element in its standard state at 25 C is zero The enthalpy of formation of H+ (aq) is also zero 0 ) ( ) (2 2 = A = A l O H H l Br H f f Calculation of H The symbol refers to the sum of Elements in their standard states may be omitted, as their enthalpies of formation are zero The coefficients of reactants and products in the balanced equation must be accounted for reactants H products H Hf f A A = A Example 6.7 Example 6.7, (Contd) Example 6.8 The First Law of Thermodynamics Thermodynamics Deals with all kinds of energy effects in all kinds of processes Two types of energy Heat (q) Work (w) The Law of Conservation of Energy Esystem = - Esurroundings The First Law E = q + w The total change in energy is equal to the sum of the heat and work transferred between the system and the surroundings Conventions q and w are positive When the heat or work enters the system from the surroundings q and w are negative When the heat or work leaves the system for the surroundings Energy Transformation and Conservation of Energy For a system or surroundings, the only possible forms of energy flow are heat, q, and work, w. The delta, A, means change in and is defined as the difference in the final and initial states. AE = q + wAE = Efinal EinitialExample Problem If 515 J of heat is added to a gas that does 218 J of work as a result, what is the change in the energy of the system? Energy Transformation and Conservation of Energy The sign resulting from the difference in the final and initial states indicates the direction of the energy flow. Negative values indicate energy is being released. Positive values indicate energy is being absorbed. Energy Transformation and Conservation of Energy First law of thermodynamics states that energy can be transformed from one form to another but cannot be created or destroyed. AEuniverse = AEsurroundings + AEsystem = 0Figure 8.10 Example 6.9 Heat Ordinarily, when a chemical reaction is carried out in the laboratory, energy is evolved as heat CH4 (g) + 2O2 (g) CO2 (g) + H2O () E = -885 kJ The combustion of methane in a Bunsen burner produces nearly 885 kJ of heat per mol The decrease in volume that takes place is a 1% work effect Work In an internal combustion engine, a significant fraction of the energy of combustion is converted to useful work The expansion of the combustion gases produces a volume and a pressure change The system does work on its surroundings Propels the car forward Overcomes friction Charges battery Like H, E is a state variable q and w are not state variables Figure 6.11 Pressure-Volume Work H and E Constant pressure Coffee-cup calorimeter H = qp Constant volume In a bomb calorimeter, there is no pressure-volume work done E = qv H and E, (Contd) H = E + PV H = E + PV The PV product is important only where gases are involved; it is negligible when only liquids or solids are involved H = E + ngRT ng is the change in the number of moles of gas as the reaction proceeds Example 8.10 96 End of Chapter 15 Fireworks are beautiful exothermic chemical reactions. 97 THERMOCHEMISTRY Engr. Edgie L. Estopace School of Chemical Engineering and Chemistry Mapua Institute of Technology