Chemistry: Atoms First Julia Burdge & Jason Overby Chapter 10 Energy Changes in Chemical Reactions Kent L. McCorkle Cosumnes River College Sacramento, CA Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Energy is the capacity to do work Radiant energy comes from the sun and is earth’s primary energy source Thermal energy is the energy associated with the random motion of atoms and molecules Chemical energy is the energy stored within the bonds of chemical substances Nuclear energy is the energy stored within the collection of neutrons and protons in the atom Potential energy is the energy available virtue of an object’s position by 6.1 Energy Changes in Chemical Reactions Heat is the transfer of thermal energy between two bodies that are at different temperatures. Temperature is a measure of the thermal energy. Temperature = Thermal Energy 400C greater thermal energy 900C Energy and Energy Changes Thermochemistry is the study of heat (the transfer of thermal energy) in chemical reactions. Heat is the transfer of thermal energy. Heat is either absorbed or released during a process. heat Surroundings 10.1 Energy and Energy Changes The system is a part of the universe that is of specific interest. The surroundings constitute the rest of the universe outside the system. System Surroundings Universe = System + Surroundings The system is usually defined as the substances involved in chemical and physical changes. Energy and Energy Changes An exothermic process occurs when heat is transferred from the system to the surroundings. “Feels hot!” 2H2(g) + O2(g) System 2H2O(l) + energy heat Surroundings Universe = System + Surroundings Energy and Energy Changes An endothermic process occurs when heat is transferred from the surroundings to the system. “Feels cold” energy + 2HgO(s) System 2Hg(l) + O2(g) heat Surroundings Universe = System + Surroundings 10.2 Introduction to Thermodynamics Thermodynamics is the study of the interconversion of heat and other kinds of energy. In thermodynamics, there are three types of systems: An open system can exchange mass and energy with the surroundings. A closed system allows the transfer of energy but not mass. An isolated system does not exchange either mass or energy with its surroundings. Introduction to Thermodynamics State functions are properties that are determined by the state of the system, regardless of how that condition was achieved. The magnitude of change depends only on the initial and final states of the system. Energy Pressure Volume Temperature The First Law of Thermodynamics The first law of thermodynamics states that energy can be converted from one form to another, but cannot be created or destroyed. ΔUsys + ΔUsurr = 0 ΔU is the change in the internal energy. “sys” and “surr” denote system and surroundings, respectively. ΔU = Uf – Ui; the difference in the energies of the initial and final states. ΔUsys = –ΔUsurr Work and Heat The overall change in the system’s internal energy is given by: ΔU = q + w q is heat q is positive for an endothermic process (heat absorbed by the system) q is negative for an exothermic process (heat released by the system) w is work w is positive for work done on the system w is negative for work done by the system Work and Heat ΔU = q + w Worked Example 10.1 Calculate the overall change in internal energy, ΔU, (in joules) for a system that absorbs 188 J of heat and does 141 J of work on its surroundings. Strategy Combine the two contributions to internal energy using ΔU = q + w and the sign conventions for q and w. Solution The system absorbs heat, so q is positive. The system does work on the surroundings, so w is negative. ΔU = q + w = 188 J + (-141 J) = 47 J Think About It Consult Table 10.1 to make sure you have used the proper sign conventions for q and w. 10.4 Calorimetry Calorimetry is the measurement of heat changes. Heat changes are measured in a device called a calorimeter. The specific heat (s) of a substance is the amount of heat required to raise the temperature of 1 g of the substance by 1°C. Specific Heat and Heat Capacity The heat capacity (C) is the amount of heat required to raise the temperature of an object by 1°C. The “object” may be a given quantity of a particular substance. heat capacity of 1 kg of water = 4.184 J 1000 g = 4184 J/C 1 g C Specific heat capacity of water Specific heat capacity has units of J/(g • °C) Heat capacity has units of J/°C heat capacity of 1 kg water Specific Heat and Heat Capacity The heat associated with a temperature change may be calculated: q = msΔT q = CΔT m is the mass. S (c) is the specific heat. ΔT is the change in temperature (ΔT = Tfinal – Tinitial). C is the heat capacity. Worked Example 10.4 Calculate the amount of heat (in kJ) required to heat 255 g of water from 25.2°C to 90.5°C. Strategy Use q = msΔT to calculate q. s = 4.184 J/g∙°C, m = 255 g, ΔT = 90.5°C – 25.2°C = 65.3°C. Solution 4.184 J 4 J or 69.7 kJ q= × 255 g × 65.3°C = 6.97×10 g∙°C Think About It Look carefully at the cancellation of units and make sure that the number of kilojoules is smaller than the number of joules. It is a common error to multiply by 1000 instead of dividing in conversions of this kind. Calorimetry Calculate the amount of heat required to heat 1.01 kg of water from 0.05°C to 35.81°C. Solution: Step 1:Use the equation q = msΔT to calculate q. q 1.01 kg 1000 g 4.184 J [35.81C 0.05C] = 151000 J = 151 kJ 1 kg g C Calorimetry A coffee-cup calorimeter may be used to measure the heat exchange for a variety of reactions at constant pressure: Heat of neutralization: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) Heat of ionization: H2O(l) → H+(aq) + OH‒(aq) Heat of fusion: H2O(s) → H2O(l) Heat of vaporization: H2O(l) → H2O(g) Calorimetry Concepts to consider for coffee-cup calorimetry: qP = nΔH System: reactants and products (the reaction) Surroundings: water in the calorimeter For an exothermic reaction: the system loses heat the surroundings gain (absorb) heat qsys = −msΔT qsurr = msΔT qsys = −qsurr The minus sign is used to keep sign conventions consistent. Worked Example 10.5 A metal pellet with a mass of 100.0 g, originally at 88.4°C, is dropped into 125 g of water originally at 25.1°C. The final temperature of both pellet and the water is 31.3°C. Calculate the heat capacity C (in J/°C) of the pellet. Strategy Water constitutes the surroundings; the pellet is the system. Use qsurr = msΔT to determine the heat absorbed by the water; then use q = CΔT to determine the heat capacity of the metal pellet. mwater = 125 g, swater = 4.184 J/g∙°C, and ΔTwater = 31.3°C – 25.1°C = 6.2°C. The heat absorbed by the water must be released by the pellet: qwater = -qpellet, mpellet = 100.0 g, and ΔTpellet = 31.3°C – 88.4°C = -57.1°C. Worked Example 10.5 (cont.) Solution 4.184 J qwater = g∙°C × 125 g × 6.2°C = 3242.6 J Thus, qpellet = -3242.6 J From q = CΔT we have -3242.6 J = Cpellet × (-57.1°C) Thus, Cpellet = 57 J/°C Think About It The units cancel properly to give appropriate units for heat capacity. Moreover, ΔTpellet is a negative number because the temperature of the pellet decreases. Constant-Volume Calorimetry Constant volume calorimetry is carried out in a device known as a constant-volume bomb. A constant-volume calorimeter is an isolated system. Bomb calorimeters are typically used to determine heats of combustion. qcal = −qrxn Constant-Volume Calorimetry To calculate qcal, the heat capacity of the calorimeter must be known. qcal = CcalΔT qrxn = −qcal qrxn = −CcalΔT Qsystem = - Qsurroundings If the system (reaction) gives off heat it is said to be exothermic this means the surroundings get warmer If the system (reaction) takes in heat it is said to be endothermic this means the surroundings get colder Bomb Calorimetry A device that enables us to measure the heat transfer Between the system and the surroundings When finding the specific heat of a metal Q of the surroundings is equal to the specific heat of the the water X the mass of the water X the change in temperature of the water Q of the metal is equal to –Q of the surroundings Then, the specific heat (c) of the metal is equal to –Q of the surroundings (mass of the metal) X (the change in temp of the metal) A piece of stainless steel of mass 25.0 g at 88.0 °C was placed in a calorimeter that contained 150.0 g of water at 20.0 °C. If the temperature of the water rose to 21.4 °C, what is the specific heat capacity of the stainless steel? The specific heat capacity of water is 4.184 J/ g·°C Ans: 0.53 J/ g ·°C A piece of graphite of mass 75.0 g at 92.0 °C was placed in a calorimeter that contained 200.0 g of water at 12.0 °C. If the temperature of the water rose to 16.8 °C, what is the specific heat capacity of graphite? The specific heat capacity of water is 4.184 J/ g·°C Ans: 0.71 J/ g·°C Calculate the heat required to raise the temperature of 50.0 g of water from 25.0 °C to 39.0 °C. The specific heat capacity of water is 4.184 J/ g·°C Ans: 2.93 kJ How much heat is absorbed by a 25.0-g sample of gold at 25.0 °C when it is immersed in boiling water? The specific heat capacity of gold is 0.128 J/ g·°C and the specific heat capacity of water is 4.184 J/ g·°C. Ans: 0.240 kJ A reaction occurs in a calorimeter that contains 120.0 g of water at 25.00 °C, and 3.60 kJ of heat is evolved. If the specific heat capacity of water is 4.184 J/ g·°C, what is the final temperature of the water? Ans: 32.17 °C Bomb Calorimetry and DH Bomb calorimetry can be used to determine the Heat of Reaction (DHrxn) the Heat of formation (DHf) and the Heat of solution (DHsoln) Q of the surroundings is equal to the total mass of the solution X specific heat of the solution X the change in temp of the solution Q of the reaction is equal to – Q of the surroundings Q of the reaction is equal to n DH. Where n is equal to the moles of reaction Therefore DH equals - (mass of soln X D temp of soln X spec. heat of soln) n of reaction When 5.00 g of ammonium nitrate are dissolved in 1.00 L of water in a calorimeter with a heat capacity of 2.30 kJ/ C, the temperature drops 0.760 C. What is the heat of solution of ammonium nitrate? Ans: +28.0 kJ/ mol If 100 mL of 1.00 M barium chloride are added to 100 mL of 1.00 M potassium sulfate in a calorimeter with a heat capacity of 1.52 kJ/ C, the temperature rises by 1.73 C. Calculate the heat of reaction for Ba2+(aq) + SO42–(aq) BaSO4(s) Ans: –26.3 kJ The standard enthalpy of formation (heat of formation) of H2O2(l) is –187.8 kJ/ mol and of H2O(l) is –285.8 kJ/ mol. (Oxygen is in it’s standard state) Calculate the reaction enthalpy for 2H2O2(l) 2H2O(l) + O2(g) Ans: –196.0 kJ Chemistry in Action: Fuel Values of Foods and Other Substances C6H12O6 (s) + 6O2 (g) 6CO2 (g) + 6H2O (l) DH = -2801 kJ/mol 1 cal = 4.184 J 1 Cal = 1000 cal = 4184 J Worked Example 10.6 A Famous Amos bite-sized chocolate chip cookie weighing 7.25 g is burned in a bomb calorimeter to determine its energy content. The heat capacity of the calorimeter is 39.97 kJ/°C. During the combustion, the temperature of the water in the calorimeter increases by 3.90°C. Calculate the energy content (in kJ/g) of the cookie. Think About It According to the label on the cookie package, a service size or 29 g, the andheat eachreleased servingby contains 150 Strategy Use qrxnis=four -Ccalcookies, ΔT to calculate the combustion of Cal. Convert the heat energy per gram to Calories per cookie servingtotodetermine verify its the cookie. Divide the released by the mass of the result.per gram C = 39.97 kJ/°C and ΔT = 3.90°C. energythe content cal 21.5 kJ 1 Cal 29 g = 1.5×102 Cal/serving × × g 4.184 kJ serving Solution qrxn = -CcalΔT = -(39.97 kJ/°C)(3.90°C) = -1.559×102 kJ Because energy content is a positive quantity, we write 1.559×102 kJ energy content per gram = = 21.5 kJ/g 7.25 g Enthalpy and Enthalpy Changes The enthalpy of reaction (ΔH) is the difference between the enthalpies of the products and the enthalpies of the reactants: ΔH = H(products) – H(reactants) Assumes reactions in the lab occur at constant pressure ΔH > 0 (positive) endothermic process ΔH < 0 (negative) exothermic process Without looking up the values, What is the standard heat of reaction for 2HN3(g) H2(g) + 3N2(g) Ans –2 DHf°[HN3(g)] Thermochemical Equations H2O(s) H2O(l) ΔH = +6.01 kJ/mol Concepts to consider: Is this a constant pressure process? What is the system? What are the surroundings? ΔH > 0 endothermic Thermochemical Equations CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH = −890.4 kJ/mol Concepts to consider: Is this a constant pressure process? What is the system? What are the surroundings? ΔH < 0 exothermic Thermochemical Equations Enthalpy is an extensive property. Extensive properties are dependent on the amount of matter involved. H2O(l) → H2O(g) Double the amount of matter 2H2O(l) → 2H2O(g) ΔH = +44 kJ/mol Double the enthalpy ΔH = +88 kJ/mol Units refer to mole of reaction as written Thermochemical Equations The following guidelines are useful when considering thermochemical equations: 1) Always specify the physical states of reactants and products because they help determine the actual enthapy changes. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) different states CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH = −802.4 kJ/mol different enthalpies ΔH = +890.4 kJ/mol Thermochemical Equations The following guidelines are useful when considering thermochemical equations: 2) When multiplying an equation by a factor (n), multiply the ΔH value by same factor. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) ΔH = − 802.4 kJ/mol 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) ΔH = − 1604.8 kJ/mol 3) Reversing an equation changes the sign but not the magnitude of ΔH. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) ΔH = − 802.4 kJ/mol CO2(g) + 2H2O(g) CH4(g) + 2O2(g) ΔH = +802.4 kJ/mol Worked Example 10.3 Given the thermochemical equation for photosynthesis, 6H2O(l) + 6CO2(g) → C6H12O6(s) + 6O2(g) ΔH = +2803 kJ/mol calculate the solar energy required to produce 75.0 g of C6H12O6. Strategy The thermochemical equation shows that for every mole of C6H12O6 produced, 2803 kJ is absorbed. We need to find out how much energy is absorbed for the production of 75.0 g of C6H12O6. We must first find out how many moles there are in 75.0 g of C6H12O6. The molar mass of C6H12O6 is 180.2 g/mol, so 75.0 g of C6H12O6 is 75.0 g C6H12O6 × 1 mol C6H12O6 180.2 g C6H12O6 = 0.416 mol C6H12O6 We will multiply the thermochemical equation, including the enthalpy change, by 0.416, in order to write the equation in terms of the appropriate amount of C6H12O6. Worked Example 10.3 (cont.) Solution (0.416 mol)[6H2O(l) + 6CO2(g) → C6H12O6(s) + 6O2(g)] and (0.416 mol)(ΔH) = (0.416 mol)(2803 kJ/mol) gives 2.50H2O(l) + 2.50CO2(g) → 0.416C6H12O6(s) + 2.50O2(g) ΔH = +1.17×103 kJ Therefore, 1.17×103 kJ of energy in the form of sunlight is consumed in the production of 75.0 g of C6H12O6. Note that the “per mole” units in ΔH are canceled when we multiply the thermochemical equation by the number of moles of C6H12O6. Think About It The specified amount of C6H12O6 is less than half a mole. Therefore, we should expect the associated enthalpy change to be less than half that specified in the thermochemical equation for the production of 1 mole of C6H12O6. 10.5 Hess’s Law Hess’s law states that the change in enthalpy for a stepwise process is the sum of the enthalpy changes for each of the steps. CH4(g) + 2O2(g) 2H2O(l) CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH = −890.4 kJ/mol 2H2O(g) ΔH = +88.0 kJ/mol CO2(g) + 2H2O(g) ΔH = −802.4 kJ/mol Enthalpy CH4(g) + 2O2(g) ΔH = −802.4 kJ ΔH = −890.4 kJ CO2(g) + 2H2O(g) CO2(g) + 2H2O(l) ΔH = +88.0 kJ Hess’s Law When applying Hess’s Law: 1) Manipulate thermochemical equations in a manner that gives the overall desired equation. 2) Remember the rules for manipulating thermochemical equations: Always specify the physical states of reactants and products because they help determine the actual enthalpy changes. When multiplying an equation by a factor (n), multiply the ΔH value by same factor. Reversing an equation changes the sign but not the magnitude of ΔH. 3) Add the ΔH for each step after proper manipulation. 4) Process is useful for calculating enthalpies that cannot be found directly. Worked Example 10.7 Given the following thermochemical equations, NO(g) + O3(g) → NO2(g) + O2(g) 3 O3(g) → O2(g) 2 O2(g) → 2O(g) ΔH = –198.9 kJ/mol ΔH = –142.3 kJ/mol ΔH = +495 kJ/mol determine the enthalpy change for the reaction NO(g) + O(g) → NO2(g) Strategy Arrange the given thermochemical equations so that they sum to the desired equation. Make the corresponding changes to the enthalpy changes, and add them to get the desired enthalpy change. Worked Example 10.7 (cont.) Solution The first equation has NO as a reactant with the correct coefficients, so we will use it as is. NO(g) + O3(g) → NO2(g) + O2(g) ΔH = –198.9 kJ/mol The second equation must be reversed so that the O3 introduced by the first equation will cancel (O3 is not part of the overall chemical equation). We also must change the sign on the corresponding ΔH value. 3 ΔH = +142.3 kJ/mol O (g) → O3(g) 2 2 These two steps sum to give: NO(g) + O3(g) → NO2(g) + O2(g) 1 3 + O2(g) O2(g) → O3(g) 2 2 1 NO(g) + O2(g) → NO2(g) 2 ΔH = –198.9 kJ/mol ΔH = +142.3 kJ/mol ΔH = –56.6 kJ/mol Worked Example 10.7 (cont.) 1 Solution We then replace 2 O2 on the left with O by incorporating the last equation. To do so, we divide the third equation by 2 and reverse its direction. As a result, we must also divide ΔH value by 2 and change its sign. 1 O(g) → O2(g) ΔH = –247.5 kJ/mol 2 Finally, we sum all the steps and add their enthalpy changes. NO(g) + O3(g) → NO2(g) + O2(g) 3 O (g) → O3(g) 2 2 1 + O(g) → O2(g) 2 NO(g) + O(g) → NO2(g) ΔH = –198.9 kJ/mol ΔH = +142.3 kJ/mol ΔH = –247.5 kJ/mol ΔH = –304 kJ/mol Think About It Double-check the cancellation of identical items–especially where fractions are involved. The standard enthalpy of formation of NO(g) is +90.3 kJ/ mol and of NO2(g) is +33.2 kJ/ mol. (Oxygen is in It’s standard state) Calculate the reaction enthalpy (heat of reaction) for 2NO(g) + O2(g) 2NO2(g) Ans: –114.2 kJ The heat of reaction for 3PbO2(s) Pb3O4(s) + O2(g) is +22.8 kJ/ mol. The standard heat of formation of Pb3O4(s) is –175.6 kJ/ mol. What is the standard heat of formation of PbO2(s)? Ans: –66.1 kJ/mol The heat of reaction for 2SO3(g) 2SO2(g) + O2(g ) is +198.4 kJ The standard heat of formation of SO2(g) is –296.8 kJ/ mol. What is the standard heat of formation of SO3(g)? Ans: –396.0 kJ/ mol Calculate the heat of reaction for H2O(l) H2(g) + ½ O2(g) from the data 2H2(g) + O2(g) 2H2O(l) Ans: DH° = –571.6 kJ +285.8 kJ For the combustion of methane, DHr = -890 kJ/mol. What is the heat of reaction for the following reaction? 4H2O(l) + 2CO2(g) 2CH4(g) + 4O2(g) Ans: +1780 kJ Calculate the heat of reaction for CH4(g) + 4S(s) CS2(l) + 2H2S(g) from the data C(s) + 2H2(g) CH4(g)DH° = –74.8 kJ C(s) + 2S(s) CS2(l) DH° = +87.9 kJ S(s) + H2(g) H2S(g) DH° = –20.6 kJ Ans: +121.5 kJ Calculate the heat of reaction for 2PbO(s) + O2(g) 2PbO2(s) from the data 2Pb(s) + O2(g) 2PbO(s) DH° = –435.8 kJ Pb(s) + O2(g) PbO2(s) DH° = –276.6 kJ Ans: –117.4 kJ Calculate the heat of reaction for C6H6(l) 3C2H2(g) from the data 2C2H2(g) + 5O2(g) 4CO2(g) + 2H2O(l) DH° = –2600 kJ 2C6H6(l) + 15O2(g) 12CO2(g) + 6H2O(l) Ans: +632 kJ DH° = –6536 kJ Calculate the heat of reaction for the combustion of methane, CH4(g), CH4(g) + 2O2(g) CO2(g) + 2H2O(l) from the data C(s) + O2(g) CO2(g) DH° = –393.5 kJ H2(g) + ½ O2(g) H2O(l) DH° = –285.8 kJ C(s) + 2H2(g) CH4(g)DH° = –74.8 kJ Ans: –890.3 kJ 10.6 Standard Enthalpies of Formation The standard enthalpy of formation (ΔH °) f is defined as the heat change that results when 1 mole of a compound is formed from its constituent elements in their standard states. C(graphite) + O2(g) CO2(g) Elements in standard states 1 mole of product ΔH f° = −393.5 kJ/mol Standard Enthalpies of Formation The standard enthalpy of formation (ΔH f°) is defined as the heat change that results when 1 mole of a compound is formed from its constituent elements in their standard states. The superscripted degree sign denotes standard conditions. 1 atm pressure for gases 1 M concentration for solutions “f” stands for formation. ΔH f° for an element in its most stable form is zero. ΔH f° for many substances are tabulated in Appendix 2 of the textbook. Standard Enthalpies of Formation The standard enthalpy of reaction (ΔH°rxn) is defined as the enthalpy of a reaction carried out under standard conditions. aA + bB → cC + dD ΔH°rxn = [cΔH °(C) + dΔH °(D) ] – [aΔH °(A) + bΔH °(B)] f f f f ΔH°rxn = ΣnΔH °(products) – ΣmΔH °(reactants) f f n and m are the stoichiometric coefficients for the reactants and products. Worked Example 10.8 Using data from Appendix 2, calculate ΔH°rxn for Ag+(aq) + Cl-(aq) → AgCl(s). Strategy Use ΔH°rxn = ΣnΔH °(products) – ΣmΔH °(reactants) and ΔH °f values f f from Appendix 2 to calculate ΔH°rxn. The ΔH °f values for Ag+(aq), Cl-(aq), and AgCl(s) are +105.9, –167.2, and –127.0 kJ/mol, respectively. Solution ΔH°rxn = ΔH f°(AgCl) – [ΔH f°(Ag+) + ΔH f°(Cl-)] = –127.0 kJ/mol – [(+105.9 kJ/mol) + (–167.2 kJ/mol)] = –127.0 kJ/mol – (–61.3 kJ/mol) = –65.7 kJ/mol Think About It Watch out for misplaced or missing minus signs. This is an easy place to lose track of them. Worked Example 10.9 Given the following information, calculate the standard enthalpy of formation of acetylene (C2H2) from its constituent elements: C(graphite) + O2(g) → CO2(g) 1 H2(g) + O2(g) → H2O(l) 2 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(l) ΔH°rxn = –393.5 kJ/mol (1) ΔH°rxn = –285.5 kJ/mol (2) ΔH°rxn = –2598.8 kJ/mol (3) Strategy Arrange the equations that are provided so that they will sum to the desired equation. This may require reversing or multiplying one or more of the equations. For any such change, the corresponding change must also be made to the ΔH°rxn value. The desired equation, corresponding to the standard enthalpy of formation of acetylene, is 2C(graphite) + H2(g) → C2H2(g) Worked Example 10.9 (cont.) Solution We multiply Equation (1) and its ΔH°rxn value by 2: 2C(graphite) + 2O2(g) → 2CO2(g) ΔH°rxn = –787.0 kJ/mol We include Equation (2) and its ΔH°rxn value as is: 1 ΔH°rxn = –285.5 kJ/mol H2(g) + 2 O2(g) → H2O(l) We reverse Equation (3) and divide it by 2 (i.e., multiply through by 1/2): 5 Think About It Remember only a=ΔH° there is rxn isΔH° f whenkJ/mol +1299.4 2CO2(g) + H2O(l) → C2H2(g)that + 2aOΔH° (g) rxn 2 just one product, just one mole produced, and all the reactants are elements in their equations standard states. Summing the resulting and the corresponding ΔH° values: rxn 2C(graphite) + 2O2(g) → 2CO2(g) ΔH°rxn = –787.0 kJ/mol ΔH°rxn = –285.5 kJ/mol 1 H2(g) + 2 O2(g) → H2O(l) 5 2CO2(g) + H2O(l) → C2H2(g) + 2 O2(g) 2C(graphite) + H2(g) → C2H2(g) ΔH°rxn = +1299.4 kJ/mol ΔH°f = +226.6 kJ/mol 10.7 Bond Enthalpy and the Stability of Covalent Molecules The bond enthalpy is the enthalpy change associated with breaking a bond in 1 mole of gaseous molecule. H2(g) → H(g) + H(g) ΔH° = 436.4 kJ/mol The enthalpy for a gas phase reaction is given by: ΔH° = ΣBE(reactants) – ΣBE(products) ΔH° = total energy input – total energy released bonds broken bonds formed Bond Enthalpy and the Stability of Covalent Molecules Bond enthalpy change in an exothermic reaction.: Bond Enthalpy and the Stability of Covalent Molecules Bond enthalpy change in an endothermic reaction: Worked Example 10.10 Use bond enthalpies from Table 10.4 to estimate the enthalpy of reaction for the combustion of methane: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) Think About It structures Use equation ΔH°rxn = ΣnΔH Strategy Draw Lewis to determine what °(products) fbonds are to–be broken and ΣmΔHare °(reactants) and data from Appendix 2 to calculate this what bonds f to be formed. enthalpy of reaction again; then compare your results using the two approaches. The difference in this case is due to two things: most tabulated bond enthalpies are averages and, by convention, we show the product of combustion as liquid water–but average bond species Bonds enthalpies to break: 4apply C–H to and 2 O=Oin the gas phase, where there is little or influence exerted neighboring molecules. Bonds no to form: 2 C=O and 4byH–O Bond enthalpies from Table 10.4: 414 kJ/mol (C–H), 498.7 kJ/mol (O=O), 799 kJ/mol (C=O in CO2), and 460 kJ/mol (H–O). Solution [4(414 kJ/mol) + 2(498.7 kJ/mol)] – [2(799 kJ/mol) + 4(460 kJ/mol)] = –785 kJ/mol