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Topic 5:Energetics 486 Version 2 2014 Table of Contents Topic 5:Energetics ......................................................................................................................................... 486 Definitions .................................................................................................................................................... 488 Temperature Conversion ........................................................................................................................ 490 Topic 5.1: Exothermic and Endothermic Reactions. ............................................................................ 490 Endothermic Reactions vs. Exothermic Reactions........................................................................... 490 Energy Diagrams ...................................................................................................................................... 491 Topic 5.2: Calculation of Enthalpy Changes........................................................................................... 496 Heat Transfer Problems ........................................................................................................................ 496 Chemistry Calorimetry Exercises ........................................................................................................ 498 Topic 5.3 Hess’s Law ................................................................................................................................... 499 Notes:Calculations using Hess's Law and Heats of Formation ..................................................... 499 Hess's Law Worksheet .......................................................................................................................... 501 Hess’s Law Challenge Worksheet ........................................................................................................ 504 Topic 5.4 Bond Enthalpies ......................................................................................................................... 506 Bond Enthalpies........................................................................................................................................ 506 Bond Enthalpies 2 .................................................................................................................................... 507 Using Hess’s Law and Bond Energy to calculate enthalpy change. ............................................... 508 Topic 5: Energetics Exam Questions ...................................................................................................... 510 Topic 5: Exam Questions Markscheme .................................................................................................. 517 Topic 15.1 Standard Enthalpy Changes of Reaction HL ...................................................................... 519 Enthalpy of Formation............................................................................................................................ 520 487 Version 2 2014 Definitions Average bond enthalpy: The average enthalpy change of breaking one mole of a bond in a gaseous atom into its constituent gaseous atoms. Born-Haber cycle: Energy cycles for the formation of ionic compounds. If there is little agreement between the theoretical and experimental values, this could indicate a degree of covalent character. Electron affinity: Enthalpy change when an electron is added to an isolated atom in the gaseous state. Endothermic: A reaction in which energy is absorbed. ΔH is +. Reactants more stable than products. Enthalpy: The internal energy stored in the reactants. Only changes in enthalpy can be measured. Entropy: A measure of the disorder of a system. Things causing entropy to increase: 1) increase of number of moles of gaseous molecules; 2) change of state from solid to liquid or liquid to gas; 3) increase of temperature Exothermic: A reaction in which energy is evolved. ΔH is –. Products more stable than reactants. Gibb’s free energy: Must be negative for reaction to be spontaneous. ΔG = ΔH – TΔS Hess’ law: Enthalpy change for a reaction depends only on difference between enthalpy of products and enthalpy of reactants. It is independent of pathway. Lattice enthalpy: The endothermic process of converting a crystalline solid into its gaseous ions, or the reverse exothermic process. The lattice enthalpy increases with decreasing size of the ions and increasing charge. Spontaneous: A reaction that has a natural tendency to occur. Standard conditions: 298 K and 1 atm. Temperature: A measure of the average kinetic energy. Standard enthalpy of vaporisation: The energy required to vaporise one mole of a liquid. Enthalpy of atomisation: The energy required to produce one mole of gaseous atoms from an element in its standard state. Bond dissociation enthalpy: The energy change when one mole of a specific bond is broken or created under standard conditions. 488 Version 2 2014 Enthalpy of Combustion: The energy released when one mole of a compound is burned in excess oxygen. Standard enthalpy of formation: The energy change when one mole of a compound is formed under standard conditions from its constituent elements in their standard states. Standard enthalpy of solution: The energy change when one mole of a substance is dissolved in an infinite amount of water under standard conditions. 489 Version 2 2014 Temperature Conversion Ko= Co + 273 Fo = 9/5 Co +32 Co = 5/9(Fo-32) Convert the following to Fahrenheit 1) 10o C ________ 2) 30o C ________ 3) 40o C ________ 4) 37o C ________ 5) 0o C ________ Convert the following to Celsius 6) 32o F ________ 7) 45o F ________ 8) 70o F ________ 9) 80o F ________ 10) 90o F ________ 11) 212o F ________ Convert the following to Kelvin 12) 0o C ________ 13) -50o C ________ 14) 90o C ________ 15) -20o C ________ Convert the following to Celsius 16) 100o K ________ 17) 200o K ________ 18) 273o K ________ 19) 350o K _________ Topic 5.1: Exothermic and Endothermic Reactions. Endothermic Reactions vs. Exothermic Reactions Classify each of the following changes as either exothermic or endothermic. Also, determine whether each change is a chemical reaction or a physical change. Explain your reasoning for each. 1. 2. 3. 4. 5. 6. An ice cube melts after being left out on the table over an extended period of time. Cracking a raw egg into a hot frying pan and cooking it. A match burns/ignites after being struck against a rough surface. The human body uses the energy provided from the digestion of food. Morning dew forming on grass and plants. Dynamite explodes in the destruction of a building. 490 Version 2 2014 Energy Diagrams 450 400 350 (d) 300 (a) Z Potential 250 Energy (kJ) (c) 200 150 (e) (f) X+Y 100 (b) 50 Reaction Coordinate 1. What is the potential energy of the products? 2. What is the potential energy of the activated complex? 3. What is the potential energy of the reactants? 4. How much activation energy is required for this reaction to occur? 5. What is the heat of reaction 6. Is the reaction exothermic or endothermic? 7. What is the activation energy of the reverse reaction? 8. What is the heat of reaction of the reverse reaction? 9. Is the reverse reaction exothermic or endothermic? 10. If a catalyst were added, which quantities, if any, would change? (give letters) 11. Would the activation energy increase, decrease, or remain unchanged? 12. Would the heat of reaction increase, decrease or remain unchanged? 491 Version 2 2014 Energy Diagrams K 900 800 700 600 X+Y (d) (b) Potential 500 Energy (kJ) (a) (e) (c) 400 Z 300 200 (f) 100 Reaction Coordinate 1. What is the potential energy of the products? 2. What is the potential energy of the activated complex? 3. What is the potential energy of the reactants? 4. How much activation energy is required for this reaction to occur? 5. What is the heat of reaction 6. Is the reaction exothermic or endothermic? 7. What is the activation energy of the reverse reaction? 8. What is the heat of reaction of the reverse reaction? 9. Is the reverse reaction exothermic or endothermic? 10. If a catalyst were added, which quantities, if any, would change? (give letters) 11. Would the activation energy increase, decrease, or remain unchanged? 12. Would the heat of reaction increase, decrease or remain unchanged? 1 492 Version 2 2014 Energy Diagrams O 90 80 70 (d) 60 (a) Z Potential 50 Energy (kJ) (c) 40 30 (e) (f) X+Y 20 (b) 10 Reaction Coordinate 1. What is the potential energy of the products? 2. What is the potential energy of the activated complex? 3. What is the potential energy of the reactants? 4. How much activation energy is required for this reaction to occur? 5. What is the heat of reaction 6. Is the reaction exothermic or endothermic? 7. What is the activation energy of the reverse reaction? 8. What is the heat of reaction of the reverse reaction? 9. Is the reverse reaction exothermic or endothermic? 10. If a catalyst were added, which quantities, if any, would change? (give letters) 11. Would the activation energy increase, decrease, or remain unchanged? 12. Would the heat of reaction increase, decrease or remain unchanged? 493 Version 2 2014 Energy Diagrams S 1800 1600 1400 1200 X+Y (d) (b) Potential 1000 Energy (kJ) (a) (e) (c) 800 Z 600 400 (f) 200 Reaction Coordinate 1. What is the potential energy of the products? 2. What is the potential energy of the activated complex? 3. What is the potential energy of the reactants? 4. How much activation energy is required for this reaction to occur? 5. What is the heat of reaction 6. Is the reaction exothermic or endothermic? 7. What is the activation energy of the reverse reaction? 8. What is the heat of reaction of the reverse reaction? 9. Is the reverse reaction exothermic or endothermic? 10. If a catalyst were added, which quantities, if any, would change? (give letters) 11. Would the activation energy increase, decrease, or remain unchanged? 12. Would the heat of reaction increase, decrease or remain unchanged? 494 Version 2 2014 Energy Diagrams W 180 160 140 (d) 120 (a) Z Potential 100 Energy (kJ) (c) 80 60 (e) (f) X+Y 40 (b) 20 Reaction Coordinate 1. What is the potential energy of the products? 2. What is the potential energy of the activated complex? 3. What is the potential energy of the reactants? 4. How much activation energy is required for this reaction to occur? 5. What is the heat of reaction 6. Is the reaction exothermic or endothermic? 7. What is the activation energy of the reverse reaction? 8. What is the heat of reaction of the reverse reaction? 9. Is the reverse reaction exothermic or endothermic? 10. If a catalyst were added, which quantities, if any, would change? (give letters) 11. Would the activation energy increase, decrease, or remain unchanged? 12. Would the heat of reaction increase, decrease or remain unchanged? 495 Version 2 2014 Energy Diagrams A 360 320 280 240 X+Y (d) (b) Potential 200 Energy (kJ) (a) (e) (c) 160 Z 120 80 (f) 40 Reaction Coordinate 1. What is the potential energy of the products? 2. What is the potential energy of the activated complex? 3. What is the potential energy of the reactants? 4. How much activation energy is required for this reaction to occur? 5. What is the heat of reaction 6. Is the reaction exothermic or endothermic? 7. What is the activation energy of the reverse reaction? 8. What is the heat of reaction of the reverse reaction? 9. Is the reverse reaction exothermic or endothermic? 10. If a catalyst were added, which quantities, if any, would change? (give letters) 11. Would the activation energy increase, decrease, or remain unchanged? 12. Would the heat of reaction increase, decrease or remain unchanged? Topic 5.2: Calculation of Enthalpy Changes Heat Transfer Problems (These questions are more physics based, but does give you a better understanding of how thermal energy moves around a system and practice the use of the specific heat equation) 1. What is the specific heat of a substance? What does it mean? 2. Surely a cotton bath mat and the tile floor in your bathroom are the same temperature. So why is it comfortable to stand on the bath mat but not on the tiled bathroom floor, which seems very cold? 3. After watching the video clip on pizza, why is it that the cheese portion of the pizza burns the roof of your mouth but the crust doesn't burn your tongue? 496 Version 2 2014 4. How many joules are required to raise the temperature of 100 g of water from 20 to 50 ºC? (12,552 J) 5. How many joules are required to raise the temperature of 100 g of iron from 20 to 50 ºC? (1350 J) 6. Why is there such a large difference in your answers between water and iron? 7. How many kilojoules are required to heat 6.3 kg of aluminum from -23 to 625 ºC ? (3633 kJ) 8. A 15.0 g block of aluminum at an initial temperature of 27.5°C absorbs 0.678 kJ of heat. What is the final temperature of the block? (The specific heat of Al is 0.902 J g-1 C°.-1) 9. A sample of aluminum absorbed 9.86 J of heat, and its temperature increased from 23.2°C to 30.5°C. What was the mass of the aluminum sample? (The specific heat of aluminum is 0.902 J g-1 C°.-1.) 10. What is the resulting temperature when 35 g of water at 75°C is mixed with 15 g of water at 15°C? (The specific heat of water is 4.184 J g-1 C°.-1.) 11. What is the resulting temperature when 100 g of water at 75°C is mixed with 30 g of ice at 0 °C? Assume that all of the ice has melted. The specific heat of water is 4.184 J g-1 C°.-1 and the heat of fusion of ice is 335 Jg-1. 12. The specific heat of lead is 0.13 Jg-1C°-1. How many J of heat would be required to raise the temperature of 15 g of lead from 22°C to 37°C? 13. How much heat is required to raise the temperature of 125 g of water from 2.2 C° to 48.6 C°? The specific heat of water is 4.184Jg-1C°-1. 14. What heat is needed to raise 3.4 kg of lead from 23oC to 58oC? Some specific heats (1.5E4 J) 15. If 23.0 kg of copper at 21.0oC absorbs 45.6 kJ of heat, what will (JoC-1kg-1) be its final temperature? (26.1oC) 16. If some aluminum at 57.0oC, cools to 24.1oC, and gives off 13.4 kJ H2O liquid 4186 of heat, what is its mass? (453 g) 2100 17. A 35.0 g of a mystery substance absorbs 314 J of heat and raises H2O ice o o its temperature by 2.14 C. What is its specific heat? (4190 J C 1 H2O steam 2010 kg-1) Heat Lost = Heat Gained Problems: Aluminum 900 1. 112 grams of a mystery liquid at 83.0 oC is mixed with 564 grams of water initially at 22 oC. The final temperature of the mixture is 33.0 oC. What is the specific heat of the mystery liquid? (Assuming no Iron 450 Copper 390 2. 3. 4. 5. heat was lost to the surroundings) (4640 J kg-1 oC-1) Lead 130 A piece of lead (c = 130 J/kgoC) at 82.0oC is mixed with 112 grams of water and an 87.5 g aluminum (c = 900. g/kgoC) calorimeter cup initially at 25.0oC. The final temperature of the system is 56.0oC. What is the mass of the piece of lead? (Assuming no heat was lost to the surroundings) (5.02 kg) 89.2 g of a mystery substance is at 99.20 oC, and it is placed in a 95.0 g iron container holding 216 ml of water both at 21.01oC. The final temperature is 23.38oC. What is the specific heat of the substance? (332 J/kg/oC) A 347 g piece of copper at 98.0oC is placed in a Styrofoam cup containing 259 ml of water at 18.0 o C. What will be the final temperature of equilibrium? (Ignore the Styrofoam) (26.9oC) A 13.5 g piece of aluminum at 93.9oC is placed in an 82.0 g iron calorimeter containing 203 g of water both at 23.0oC. What will be the final temperature? (24.0oC) 497 Version 2 2014 6. If you drop a 16 g ice cube at 0.0 oC into a Styrofoam cup containing 241 ml of water at 20.0 oC what will be the final temperature? (13.8oC) 7. You take an ice cube out of the freezer at -17.0oC, and drop it into a 67.0 g aluminum cup containing 308 g of water at 23.0 oC. The final temperature is observed to be 12.7oC. What is the mass of the ice cube? (33.0 g) Chemistry Calorimetry Exercises 1. When 12.29 g of finely divided brass (60% Cu, 40% Zn) at 95.0oC is quickly stirred into 40.00 g of water at 22.0oC in a calorimeter, the water temperature rises to 24.0oC. Find the specific heat of brass. Hints: • The heat lost by the brass is gained by the surroundings (the water plus the calorimeter). What relation can you therefore write between qbrass and qsurr? • Since no information is given about the heat capacity of the calorimeter, you should assume it is negligible. • The final temperature of the brass is the same as the final temperature of the water. The specific heat of water, s(H2O), is 4.184 J g-1 oC-1. (Answers: qsurr = q(H2O) = 334.7 J; qbrass = - 334.7 J; sbrass = 0.38 J g-1 oC-1). 2. In an experiment, 400. mL of 0.600 M HNO3(aq) is mixed with 400. mL of 0.300 M Ba(OH)2(aq) in a constant-pressure calorimeter having a heat capacity of 387 J/oC. The initial temperature of both solutions is the same at 18.88oC, and the final temperature of the mixed solution is 22.49oC. Calculate the heat of neutralization in kJ per mole of HNO3. Hints: The heat evolved in the neutralization reaction is gained by the surroundings (the mixed solution plus the calorimeter). What relation can you therefore write between qrxn and qsurr? There are two contributions to qsurr. What are they? What assumptions (if any) need to be made in calculating these contributions? Is this a limiting reagent problem, or are reactants supplied in the stoichiometric ratio given by the equation? (Why do we care about this?) We want our answer in kJ per mole of HNO3. How do we calculate that? (Answers: ΔT = 3.61oC; qsolution = 12083.4 J; qcalorimeter = 1397.1 J; qsurr = 13481 J; qrxn = 13481 J; ΔHneut(kJ/mol HNO3) = - 56.2 kJ/ mol HNO3). 3. The reaction of 1.0 mol of C to form carbon monoxide in the reaction 2 C(s) + O2(g) 2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of C according the above information? (940 kJ) 4. When 1.0 mole of ethanol (C2H5OH) is burned, it releases -1367 KJ. How much heat is released if 50 g of ethanol burn? (1500 kJ) 5. Octane (C8H18) is a component in gasoline used to fuel a car. If the heat of combustion of Octane is -5471.5 kJ/mol of octane, how much energy is provided by burning 10.0 gallons? The density of octane is 0.7028 g/mL. (1.28x106 kJ) 6. 26 g of water at 18oC are mixed with 49 g of water at 70oC. Find the final temperature of the system. (52 oC) 7. 84 g of water at 22oC are mixed with 150 g of ethanol (C = 2.44 J/gC) at 88oC. Find the final temperature of the system. (55.7 oC) 8. 75.0 g of water at 30oC are mixed with 83.8 g of a solid metal at 600oC. The final temperature of the system is 50oC. What is the metal? (W) 498 Version 2 2014 9. When 1.00 dm3 of 1.00 mol dm-3 Ba(NO3)2 solution at 25.0C is mixed with 1.00 dm3 of 1.00 mol dm-3 Na2SO4 solution at 25.0C in a calorimeter, the white solid BaSO4 forms and the temperature of the mixture increases to 28.1C. Assuming that the calorimeter absorbs only a negligible quantity of heat, that the specific heat capacity of the solution is 4.18 J/C-g, and that the density of the final solution is 1.0 g/cm3, calculate the enthalpy changer per mole of BaSO4 formed. (25.9 kJ/mol) 10. In a coffee-cup calorimeter, 1.60 g of NH4NO3 is mixed with 75.0 g of water at an initial temperature of 25.00C. After dissolution of the salt, the final temperature of the calorimeter contents is 23.34C. Assuming the solution has a heat capacity of 4.18 J/C-g and assuming no heat loss to the calorimeter, calculate the enthalpy change for the dissolution of NH4NO3 in units of kJ/mol. (26.0 kJ/mol) Phase Diagram Worksheet Refer to the phase diagram below when answering the questions on this worksheet: 1. What is the normal freezing point of this substance? 2. What is the normal boiling point of this substance? 3. What is the normal freezing point of this substance? 4. If I had a quantity of this substance at a pressure of 1.25 atm and a temperature of 3000 C and lowered the pressure to 0.25 atm, what phase transition(s) would occur? 5. At what temperature do the gas and liquid phases become indistinguishable from each other. 6. If I had a quantity of this substance at a pressure of 0.75 atm and a temperature of -1000 C, what phase change(s) would occur if I increased the temperature to 6000 C? At what temperature(s) would they occur? Topic 5.3 Hess’s Law Notes:Calculations using Hess's Law and Heats of Formation Enthalpy of reaction values have been determined experimentally for numerous reactions, and these ∆H values may be used to calculate ∆H values for other reactions involving the same chemical species. The reason 499 Version 2 2014 this is possible is that enthalpy H is a state property so ∆H is independent of path. (Similarly, the height of a mountain above sea level is independent of the path you follow to climb the mountain.) Because ∆H is independent of path, we can determine the enthalpy of foods by burning them in a bomb calorimeter in the laboratory to produce the same products that are obtained by the complicated metabolic pathways in our body! There are two principle methods used to calculate ∆H values for a reaction, both of which are based on the idea that ∆H for a reaction is independent of the path used to go from reactants to products. The first makes o use of Hess's Law while the second employs tabulated heats of formation H f (kJ/mol). Use of Hess's Law to Calculate ∆H Hess's Law states that ∆H for a reaction can be found indirectly by summing ∆H values for any set of reactions which sum to the desired reaction. Usually before reactions are added together, some of them must be reversed and/or multiplied by a factor n in order that they sum to the desired reaction. In this process the rules are: • Whenever you multiply a reaction by n, ∆H for the reaction is also multiplied by n. • If you reverse a reaction, ∆H changes sign. Problem (1) below is an example of how this procedure is used. Use of Tabulated Heats of Formation to Calculate ∆H o The Standard Heat of Formation H f (kJ/mol) for a compound is the heat absorbed (or released) in forming one mole of the compound from its elements in their standard states at 1 bar o (≈ 1 atm) pressure and the specified temperature (usually 25 oC). Thus H f (kJ/mol) for acetone CH3COCH3 is the heat of the reaction 3 C(s) + 3 H2(g) + 1/2 O2(g) CH3COCH3(l) ∆H = – 246.8 kJ o o and so H f (CH3COCH3(l)) = – 246.8 kJ/mol). By definition, H f (kJ/mol) = 0 for any element in its standard state at 25oC and 1 bar. These tables may be used to calculate the standard enthalpy change, o , for any reaction for which the Hrxn heats of formation of all reactants and products are known: o Hrxn n products o H f ( prod) prod n react o H f (react) reactants This equation tells us to sum the enthalpies of formation of each product multiplied by its stoichiometric coefficient in the reaction equation and then to subtract the enthalpy of formation of each reactant multiplied by its stoichiometric coefficient. We use this equation to work problem (2). o applies for a balanced equation with specific stoichiometric amounts. If a different number of Hrxn moles reacts, the heat absorbed or evolved will change proportionately (problem 3). 1. From the following heats of reaction 2 SO2(g) + O2(g) 2 SO3(g) ∆H = – 196 kJ (a) 2 S(s) + 3 O2(g) 2 SO3 (g) ∆H = – 790 kJ (b) ∆H = ? kJ (c) calculate the heat of reaction for S(s) + O2(g) SO2(g) Method: Use Hess’s Law to solve this problem: 500 Version 2 2014 • Identify a species in the target equation (c) which is on the correct side in only one of the listed equations, (a) or (b). Multiply the entire equation, and its ∆H value by the factor n necessary to make the stoichiometric coefficient for the species identical to that in equation (c). • Reverse a listed equation, (a) or (b), and change the sign of its ∆H value if it contains a species which is on the wrong side of the target equation (c); next multiply the entire reversed equation by the factor n necessary to make the stoichiometric coefficient for the species identical to that in equation (c). The ∆H value for the rewritten equation is ( - n) times that of the original equation. (Ignore any species present in both equations (a) and (b).) • Test to see if your rewritten equations now sum to the desired equation (c). If they do, the ∆H value for equation (c) is the sum of the ∆H values of the rewritten equations. Answer: (1/2) Eq(b) + ( - 1/2) Eq(a) = Eq(c); Thus, (1/2)H (b) + ( – 1/2)H (a) = H (c) so H(c) =(1/2)( – 790 kJ) + ( – 1/2)( – 196 kJ) = kJ 2. Calculate the standard reaction enthalpy for the photosynthesis reaction, o Hrxn = ? kJ 6 CO2(g) + 6 H2O(l) C6H12O6(s) + 6 O2(g) Note: Search for the heat of formation of glucose, C6H12O6(s) o o o Answer: Use Eq (2) with H f (CO2(g)) = – 393.5 kJ/mol, H f (H2O(l)) = – 285.8 kJ/mol, H f (C6H12O6(s)) = – 1274.5 o kJ/mol, and H f (O2(g)) = 0 kJ/mol. Thus Ho = 2801.3 kJ ≈ 2801 kJ for the photosynthesis reaction. 3. Is the photosynthesis reaction above endothermic or exothermic? How much heat is absorbed or evolved if 11.0 g of CO2(g) reacts completely with excess water to form glucose and oxygen gas? Answers: Endothermic. 117 kJ of heat is absorbed. Hess's Law Worksheet For each problem below, write the equation and show your work. Always show units. 1. Consider the following hypothetical reactions: A B ∆H = +30 kJ B C ∆H = +60 kJ Use Hess’s law to calculate the enthalpy change for reaction A C. Construct an enthalpy diagram for substances A, B, and C, and show how Hess’s law applies. 2. Suppose you are given the following hypothetical reactions: XY ∆H = -40 kJ X Z ∆H = -95 kJ Use Hess’s law to calculate the enthalpy change for the reaction Y Z. Construct an enthalpy diagram for substances X, Y, and Z, and discuss its relevance to Hess’s law. 3. From the following heats of reaction: 2H2 (g) + O2 (g) 2H2O (g) 3O2 (g) 2O3 (g) Calculate the heat of the reaction 3H2 (g) + O3 (g) 3H2O (g) ∆H = -483.6 kJ ∆H = +284.6 kJ 4. From the following enthalpies of reaction: H2 (g) + F2 (g) 2HF (g) ∆H = - 537kJ 501 Version 2 2014 C (s) + 2 F2 (g) CF4 (g) 2C (s) + 2 H2 (g) C2H4 (g) ∆H = - 680 kJ ∆H = + 52.3 kJ Calculate the ΔH for the reaction of ethylene with F2. C2H4 (g) + 6F2 (g) 2CF4 (g) + 4HF(g) 5. Given the following data: N2 (g) + O2 (g) 2NO (g) 2NO (g) + O2 (g) 2NO2 (g) 2N2O (g) 2N2 (g) + O2 (g) ∆H = + 180.7 kJ ∆H = - 113.1 kJ ∆H = - 162.3 kJ Use Hess’s law to calculate ΔH for the reaction. N2O (g) + NO2 (g) 3NO (g) (Answers: 1a) 90 kJ 2a) -55 kJ3) -867.7 kJ4)-2486.3 kJ5) 156.1 kJ) 6. Calculate ∆H for the reaction: C2H4 (g) + H2 (g) C2H6 (g), from the following data. C2H4 (g) + 3O2 (g) 2 CO2 (g) + 2H2O(l) ∆H = -1411. kJ C2H6(g) + 3½O2 (g) 2CO2 (g) + 3H2O(l) ∆H = -1560. kJ H2 (g) + ½O2 (g) H2O(l) ∆H = -285.8 kJ 7. Calculate ∆H for the reaction 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g), from the following data. N2 (g) + O2 (g) 2NO (g) N2 (g) + 3H2 (g) 2NH3 (g) 2H2 (g) + O2 (g) 2H2O (g) ∆H = -180.5 kJ ∆H = -91.8 kJ ∆H = -483.6 kJ 8. Find ∆H for the reaction 2H2(g) + 2C(s) + O2(g) C2H5OH(l), using the following thermochemical data. C2H5OH (l) + 2O2 (g) 2CO2 (g) + 2H2O (l) C (s) + O2 (g) CO2 (g) H2 (g) + ½O2 (g) H2O (l) ∆H = -875. kJ ∆H = -394.51kJ ∆H = -285.8 kJ 9. Calculate ΔH for the reaction CH4 (g) + NH3 (g) HCN (g) + 3 H2 (g), given: N2 (g) + H2 (g) 2NH3 (g) C (s) + 2H2 (g) CH4 (g) H2 (g) + 2C (s) + N2 (g) 2HCN (g) ∆H = -91.8 kJ ∆H = -74.9 kJ ∆H = +270.3 kJ 10. Calculate ∆H for the reaction 2 Al (s) + 3 Cl2 (g) 2 AlCl3 (s) from the data. 2Al (s) + 6HCl (aq) 2AlCl3 (aq) + 3H2 (g) ∆H = -1049. kJ HCl (g) HCl (aq) ∆H = -74.8 kJ 502 Version 2 2014 H2 (g) + Cl2 (g) 2 HCl (g) AlCl3 (s) AlCl3 (aq) ∆H = -1845. kJ ∆H = -323. kJ 11. Calculate ΔHº for Mg(s) + ½ O2(g) MgO(s) given the equations: Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g) ΔHº = –462 kJ MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l) ΔHº = –146 kJ 2H2(g) + O2(g) 2H2O(l) ΔHº = –572 kJ Then draw the Enthalpy Diagram for the overall reaction. 12. Calculate ΔHº for C2H4(g) + 2H2(g) 2CH4(g) given the equations: CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔHº = –890 kJ C2H4(g) + 3O2(g) 2CO2(g) + 2H2O(l) ΔHº = –1411 kJ 2H2(g) + O2(g) 2H2O(l) ΔHº = –572 kJ Then draw the Enthalpy Diagram for the overall reaction. ΔH = -602 kJ ΔH = -203 kJ Calculate ΔH (kJ) for the reaction: NH3(g) + 5/2 O2(g) 2 NO(g) + 3 H2O(g) given the following: ½ N2(g) + ½ O2(g) NO(g) ΔH= 90.4 kJ mol-1 ½ N2(g) + 3/2 H2(g) NH3(g) ΔH= -46.0 kJ mol-1 H2(g) + ½ O2 (g) H2O(g) ΔH= -241.8 kJ mol-1 ΔH = -452.6 kJ 13. Calculate ΔH (kJ) for the reaction: Ca2+(aq) + 2OH-(aq) + CO2(g) CaCO3(s) + H2O(l) given the following: CaCO3(s) CaO(s) + CO2(g) ΔH = 178.1 kJ mol-1 CaO(s) + H2O(l) Ca(OH)2(s) ΔH = -64.8 kJ mol-1 Ca(OH)2(s) Ca2+(aq) + 2OH-(aq) ΔH = -11.7 kJ mol-1 ΔH = -101.6 kJ 14. Calculate ΔH (kJ) in kJ for the reaction: 2C(s) + H2(g) C2H2(g) given the following: C2H2(g) + 5/2O2 (g) 2CO2(g) + H2O(l) ΔH= - 1300 kJ mol-1 C(s) + O2(g) CO2(g) ΔH= -394 kJ mol-1 H2(g) + ½ O2(g) H2O(l) ΔH= -286 kJ mol-1 ΔH = 226 kJ 15. The bombardier beetle uses an explosive discharge as a defensive mechanism. The chemical reaction involved is the oxidation of hydroquinone by hydrogen peroxide to produce quinone and water. C6H4(OH)2(aq) + H2O2(aq) C6H4O2(aq) + 2H2O(l) Calculate ΔH in kJ for the reaction given the following: C6H4(OH)2(aq) C6H4O2(aq) + H2(g) H2(g) + O2(g) H2O2(aq) H2(g) + ½ O2(g) H2O(g) H2O(g) H2O(l) ΔH = +177.4 kJ mol-1 ΔH = -191.2 kJ mol-1 ΔH= -241.8 kJ mol-1 ΔH = -43.8 kJ mol-1 16. The enthalpies of these reactions at 25°C are given in kJ. C4H9OH (l) + 6 O2(g) (C 2H5)2O(l) + 6 O2(g) 4 CO2 (g) + 5 H2O (g) 4 CO2 (g) + 5 H2O (g) 503 ΔH = -202 kJ ∆H° in kJ mol-1 -2456 - 2510 Version 2 2014 From a consideration of the above information, calculate the change in enthalpy for C4H9OH (l) (C 2H5)2O (l) 17. The enthalpies of three reactions at 25°C are given in kJ. ∆H° in kJ mol-1 2 H2 (g) + O2 (g) 2 H2O (l) -572 2 C2H6 (g) + 7 O2 (g) 4 CO2(g) + 6 H2O(l) - 3120 C2H4 (g) + 3 O2 (g) 2 CO2 (g) + 2 H2O (l) - 1411 Use this information to calculate the change in enthalpy for C2H6 (g) C2H4 (g) + H2 (g) Is the reaction endothermic or exothermic? Explain. 18. The enthalpies of three reactions at 25°C are given in kJ. ∆H° in kJ mol-1 Mg (s) + 1/2 O2(g) MgO (s) -601.7 C(s) + O2 (g) CO2 (g) -393.5 Mg (s) + C (s) + 3/2 O2 (g) Mg CO3 (g) -1095.8 Use this information to calculate the change in enthalpy for MgO (s) + CO2 (g) MgCO3 (s) 19. The enthalpies of three reactions at 25°C are given in kJ. C(s) + 1/2 O2 (g) CO (g) C(s) + O2 (g) CO2 (g) CO (g) + 1/2 O2 (g) CO2 (g) H2 (g) + 1/2 O2 (g) H2O (g) Use this information to calculate the change in enthalpy for C(s) + H2O (g) CO (g) + H2 ∆H° in kJ mol-1 - 110.5 - 393.5 -283.0 -241.8 (g) 20. The enthalpies of three reactions at 25°C are given in kJ. N2 (g) + 2O2 (g) 2 NO2 (g) 2 NO (g) + O2 (g) 2 NO2 (g) ∆H° in kJ mol-1 + 33.2 - 57.1 From a consideration of the above information, calculate the change in enthalpy for N2 (g) + O2 (g) 2 NO (g) Hess’s Law Challenge Worksheet 1. Use the data below, as appropriate, to answer these questions. Cl2(g) 2 Cl(g) H2(g 2 H(g) HCl(g) H(g) + Cl(g) Hz(g) + 1/2 O2(g) HzO(g) O2(g) 2 O(g) ∆H= +242kJmol-1 ∆H = + 436 kJ mol-1 ∆H = + 431 kJ mol-1 ∆H = -242kJmol'1 ∆H = + 496 kJ mol-1 a) State Hess's Law. 504 Version 2 2014 b) i. Calculate the standard enthalpy change for the following reaction. H2(g) + Cl2(g) 2HCl(g) ii. ln tne dark, the reaction between hydrogen and chlorine starts when the mixture is heated to above 600'C. Suggest why heat is necessary in order to start the reaction. iii. What happens in the first step in the reaction between hydrogen and chlorine? Give a reason why this occurs first. c) Calculate a value for the mean bond enthalpy of the H-O bond in water. 2. a) Define the term standard enthalpy of formation (∆Hf). b) Give the chemical equation for which he enthalpy change is the standard enthalpy of formation of methane. c) Use the following data to calculate a value for the enthalpy of formation of methane C(s) + O2(g) CO2(g) ∆H = -394 kJ mol-1 H2(g) + ½ O2(g) H2O(g) ∆H = -242kJmol-1 CH4(g) + 2O2(g) CO2(g) + 2H2O(g) ∆H = - 802 kJ mol-1 d) Use the following data to calculate the mean bond enthalpy values for the C-H and C-C bonds. CH4(g) C(g) + 4H(g) ∆H = + 1648 kJ mol-1 C2H6(g) 2C(g) + 6H(g) ∆H = + 2820 kJ mol-1 3) a) Write a chemical equation, including state symbols, for the reaction which is used to define the enthalpy of formation of aluminium chloride (AlCl3). b) State the additional conditions necessary if the enthalpy change for this reaction is to be the standard enthalpy of formation, Hf at 298 K. c) Use the standard enthalpies of formation given below to calculate a value for the standard enthalpy change for the following reaction. AlCl3(s) + 6 H2O(l) AlCl3.6H2O(s) 4) a) Explain the meaning of the term standard enthalpy of combustion of a compound. b) Write an equation to represent the enthalpy of combustion of methanol. c) The apparatus shown was used to find the enthalpy of combustion of methanol. In this experiment, 0.75 grams of methanol was burned. The water in the apparatus, which had a volume of 100cm3, rose in temperature from 18.5oC to 50.3oC. i. Calculate the heat released by the 0.75 grams of methanol, given that the specific heat capacity of water is 4.2Jg-1 K-1. ii. Calculate the enthalpy of combustion of methanol iii. The actual value of the enthalpy of combustion of methanol is -726 kJ mol-l. Account for the difference between this value and the experimental value. 505 Version 2 2014 Topic 5.4 Bond Enthalpies Bond Enthalpies The energy required to break a covalent bond in a gaseous substance is called the bond energy for the chemical bond. For example, for the diatomic molecule HCl is defined as the standard enthalpy change for the reaction: HCl (g) H (g) + Cl (g) ∆H° = 431 kJ mol-1 The enthalpy change for the equation H2O (g) O (g) + 2 H (g) is ∆Hreaction = + 925 kJ. Since water is made up of two O-H covalent bonds, the value of ∆Hreaction is the energy required to break these bonds. Since ∆Hreaction is positive, at constant pressure 925 kJ of energy is required to break two moles H-O covalent bonds in 1 mole of H2O. The average molar bond enthalpy for a hydrogen to oxygen bond is equal to one half of the total energy input as heat required to break two moles of hydrogen to oxygen bonds or about 463 kJ mol-1. It has been found that the hydrogen to oxygen molar bond enthalpy in a variety of compounds is about the same as the H-O bond enthalpy in water. Bond enthalpies are a function of the particular atoms forming the bond rather than the particular substance in which they are incorporated. Some typical molar bond enthalpies are listed below: Bond Bond enthalpy Bond Bond enthalpy Bond Bond enthalpy O-H O-O C-O O=O C=O C-C C=C C= C C-H 464 142 351 502 730 347 615 811 414 C-F C-Cl C-Br C-N C=N C= N N-H N-N N=N 439 439 276 293 615 890 390 159 418 N= N F-F Cl-Cl Br-Br H-H H-F H-Cl H-Br H-S 945 155 243 192 435 565 431 368 364 In general it is helpful to picture a chemical reaction as taking place in two steps. First the chemical bonds in the reactants are broken down into their constituent atoms and then the atoms are rejoined to form the product molecules. The first step requires an input of energy to break the chemical bonds of the reactants. The second step releases energy as new chemical bonds are formed. The enthalpy change for the reaction, ∆Hrxn, is the difference between the energy required to break the chemical bonds of the reactants and the energy released in the formation of the chemical bonds for the products. Example 1: Calculate the enthalpy of reaction for the following reaction: CH4 (g) + Cl2 (g) CH3Cl (g) + HCl(g) In the above reaction it is necessary to break one C-H bond and one Cl-Cl bond and to form one C-Cl bond and one H-Cl bond. Reactants: (Bonds broken) 4 C-H 414 kJ x 4 = 1656 kJ 1 Cl-Cl 243 kJ x 1 = 243 kJ Total = 1899 kJ (absorbed ∆H is +) Products (Bonds formed) 3 C-H 414 kJ x 3 = 1242 kJ 1 C-Cl 439 kJ x 1 = 439 1 H-Cl 431 kJ x 1 = 431 Total = 2112 kJ (Released ∆H is -) 506 Version 2 2014 Change in enthalpy = - 213 kJ Exercises: 1. The enthalpy change for the reaction: ClF3 (g) Cl (g) + 3 F (g) is 514 kJ. Using this information and the bond enthalpies on the previous page, calculate the average Cl-F bond energy in ClF3. 2. Use the bond enthalpies on the previous page to calculate ∆Hreaction for the following reaction: CH4 (g) + 2 Cl2 (g) CH2Cl2 (g) + 2 HCl (g) 3. Hydrazine, N2H4 and its derivatives are used as rocket fuels. Use the molar bond enthalpies on the previous page to calculate the molar enthalpy of formation of N2H4 (g). 4. The formation of water from oxygen and hydrogen involves the reaction: 2 H2 (g) + O2 (g) 2 H2O (g) Use the bond enthalpies on the previous page to calculate the enthalpy of formation of H 2O (g). 5. The formation of ammonia from nitrogen and hydrogen involves the reaction: 3 H2 (g) + N2 (g) 2 NH3 (g) Use the bond enthalpies on the previous page to calculate the enthalpy of formation of NH 3 (g) . Bond Enthalpies 2 1. 2. 3. 4. Which process releases energy: breaking a bond or forming a bond? Which process requires energy: breaking a bond or forming a bond? Define bond energy. If the energy used to break bonds is greater than the energy released in the formation of new bonds, is the reaction endothermic or exothermic? 5. If the energy used to break bonds is less than the energy released in the formation of new bonds, is the reaction endothermic or exothermic? 6. Look at the table of selected bond energies to the right. a. Which is the strongest single bond on this table? b. Which is the weakest single bond on this table? 7. (Note: This is not a question. It is an explanation of how to calculate the total energy from a reaction, based on bond energies. There are supposed to be empty spaces in the table.) Let’s examine the electrolysis of water (fig 4 on page 587). The general reaction is 2H2O → 2H2 + O2 or H−O−H + H−O−H → H−H + H−H + O=O The overall heat of reaction can be calculated as follows: Thus, this is an endothermic reaction (energy required) that absorbs 466 kJ. 507 Version 2 2014 The thermochemical equation is 2H2O + 466 kJ → 2H2 + O2 Calculate the heat of reaction for H2 + Cl2 → 2HCl (by completing a table similar to the one above. Write the thermochemical equation for this reaction. 8. Calculate the theoretical value molar heat of reaction for burning paraffin (C25H52) in oxygen. a. Draw Lewis structures for O2, H2O, and CO2. b. Write the balanced equation for the combustion of C25H52. c. Fill in this table to calculate the theoretical molar heat of combustion for C25H52 9. Calculate the molar heat of combustion & the specific heat of combustion C2H5OH (shown right). for Using Hess’s Law and Bond Energy to calculate enthalpy change. 1. Given the enthalpy changes for the reactions below 2 H2O2 (aq) 2 H2O (l) + O2 (g) ∆H = -200 kJ mol-1 2 H2 (g) + O2 (g) 2 H2O (l) ∆H = -600kJ mol-1 what will be the enthalpy change for : H2 (g) + O2 (g) H2O2 (aq) ? A. -200 kJ mol-1 B. -400 kJ mol-1 C. -600 kJ mol-1 D. -800 kJ mol-1 2. Iron and chlorine react directly to form iron(III) chloride, not iron(II) chloride. Therefore it is not possible to directly measure the enthalpy change for the reaction Fe(s) + Cl2 (g) FeCl2 (s) The enthalpy changes for the formation of iron(III) chloride from the reaction of chlorine with iron and with iron(II) chloride are given below. Use these to calculate the enthalpy change for the reaction of iron with chlorine to form iron (II) chloride. 2 Fe (s) + 3 Cl2 (g) 2 FeCl3 (s) ΔH = -800 kJ mol-1 2 FeCl2 (s) + Cl2 (g) 2 FeCl3 (s) ΔH = -120 kJ mol-1 3. The Romans used calcium oxide (CaO) as mortar in stone structures. The CaO was mixed with water to give Ca(OH)2, which slowly reacted with CO2 in the air to give limestone. Ca(OH)2 (s) + CO2 (g) CaCO3 (s) + H2O (l) Use the enthalpies of the following reactions to calculate the enthalpy change for the reaction above. Ca (s) + O2 (g) + H2 (g) Ca(OH)2 (s) ΔH = -986.1 kJ/mol C (s) + O2 (g) CO2 (g) ΔH = -394.1 kJ/mol H2(g) + 1/2 O2(g) H2O(l) ΔH = -285.8 kJ/mol Ca (s) + C (s) + 3/2 O2 (g) CaCO3(s) ΔH = -1206.9 kJ/mol 508 Version 2 2014 4. Which of the following equations is equivalent to the bond enthalpy of the carbon-oxygen bond in carbon monoxide? A. CO (g) C(g) + O (g) B. CO (g) C(s) + O (g) C. CO (g) C(s) + ½ O2 (g) D. CO (g) C(g) + ½ O2 (g) 5. The bond enthalpy of the N-O bond in nitrogen dioxide is 305 kJ mol-1. If that of the bonds in the oxygen molecule and the nitrogen molecule are 496 kJ/mol and 944 kJ/mol respectively, what will be the enthalpy change for the following reaction? N2 (g) + 2 O2 (g) 2 NO2 (g) A. 716 kJ mol-1 B. 1135 kJ mol-1 C. 1326 kJ mol-1 D. 1631 kJ mol-1 6. The enthalpy change for the following reaction is +688 kJ mol-1. N2 (g) + 3 Cl2 (g) 2 NCl3 (g) Calculate the bond enthalpy of the N-Cl bond, given that the bond enthalpies of the nitrogen molecule and the chlorine molecule are 944 kJ mol-1 and 242 kJ mol-1 respectively. 7. Use bond energy data to calculate the enthalpy change when cyclopropane reacts with hydrogen to form propane. The actual value is –159 kJ mol-1. Give the reasons why you think this differs from the value you have calculated. [Bond energies in kJ mol-1: C-C = 348; C-H = 412; H-H = 436] The reaction is shown in the following chemical equation: 509 Version 2 2014 Topic 5: Energetics Exam Questions 1. When some solid barium hydroxide and solid ammonium thiosulfate were reacted together, the temperature of the surroundings was observed to decrease from 15 ºC to – 4 ºC. What can be deduced from this observation? A. B. C. D. The reaction is exothermic and ∆H is negative. The reaction is exothermic and ∆H is positive. The reaction is endothermic and ∆H is negative. The reaction is endothermic and ∆H is positive. (Total 1 mark) 2. Which process represents the C–Cl bond enthalpy in tetrachloromethane? A. B. C. D. CCl4(g) → C(g) + 4Cl(g) CCl4(g) → CCl3(g) + Cl(g) CCl4(l) → C(g) + 4Cl(g) CCl4(l) → C(s) + 2Cl2(g) (Total 1 mark) 3. Some water is heated using the heat produced by the combustion of magnesium metal. Which values are needed to calculate the enthalpy change of reaction? I. The mass of magnesium II. The mass of the water III. The change in temperature of the water A. B. C. D. I and II only I and III only II and III only I, II and III (Total 1 mark) 3. In a reaction that occurs in 50 g of aqueous solution, the temperature of the reaction mixture increases by 20 °C. If 0.10 mol of the limiting reagent is consumed, what is the enthalpy change (in kJ mol–1) for the reaction? Assume the specific heat capacity of the solution 4. = 4.2 kJ kg–1 K–1. A. B. C. D. –0.10 × 50 × 4.2 × 20 –0.10 × 0.050 × 4.2 × 20 50 4.2 20 0.10 0.050 4.2 20 0.10 (Total 1 mark) 5. Use the average bond enthalpies below to calculate the enthalpy change, in kJ, for the following reaction. H2(g) + I2(g) → 2HI(g) Bond Bond energy / kJ mol–1 H–H 440 I–I 150 510 Version 2 2014 H–I A. B. C. D. 300 +290 +10 –10 –290 (Total 1 mark) 6. Which substance does not conduct electricity? A. B. C. D. Solid zinc Molten zinc Solid zinc chloride Molten zinc chloride (Total 1 mark) 7. 1.0 g of sodium hydroxide, NaOH, was added to 99.0 g of water. The temperature of the solution increased from 18.0 ºC to 20.5 ºC. The specific heat capacity of the solution is 4.18 J g–1 K–1. Which expression gives the heat evolved in kJ mol–1? A. 2.5 100.0 4.18 1000 40.0 B. 2.5 100.0 4.18 1000 40.0 C. 2.5 100.0 4.18 40.0 1000 D. 2.5 1.0 4.18 40.0 1000 (Total 1 mark) 8. Which is true for a chemical reaction in which the products have a higher enthalpy than the reactants? Reaction ∆H A. endothermic positive B. endothermic negative C. exothermic positive D. exothermic negative (Total 1 mark) 9. Consider the reaction between magnesium and hydrochloric acid. Which factors will affect the reaction rate? I. The collision frequency of the reactant particles II. The number of reactant particles with E ≥ Ea III. The number of reactant particles that collide with the appropriate geometry A. B. C. D. I and II only I and III only II and III only I, II and III (Total 1 mark) 511 Version 2 2014 10. What is the energy, in kJ, released when 1.00 mol of carbon monoxide is burned according to the following equation? 2CO(g) + O2(g) → 2CO2(g) A. B. C. D. ΔHo = –564 kJ 141 282 564 1128 (Total 1 mark) 11. Which of the following reactions are exothermic? A. B. C. D. I. CH4 + 2O2 → CO2 + 2H2O II. NaOH + HCl → NaCl + H2O III. Br2 → 2Br I and II only I and III only II and III only I, II and III (Total 1 mark) 12. The specific heat of iron is 0.450 J g–1 K–1. What is the energy, in J, needed to increase the temperature of 50.0 g of iron by 20.0 K? A. 9.00 B. 22.5 C. 45.0 D. 450 (Total 1 mark) 13. Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of hydrogenation of ethene to form ethane. C2H4(g) + H2(g) → C2H6(g) John used the average bond enthalpies from Table 10. Marit used the values of enthalpies of combustion from Table 12. (a) Calculate the value for the enthalpy of hydrogenation of ethene obtained using the average bond enthalpies given in Table 10. ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... (2) (b) Marit arranged the values she found in Table 12 into an energy cycle. 512 Version 2 2014 Calculate the value for the enthalpy of hydrogenation of ethene from the energy cycle. ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... (1) (c) Suggest one reason why John’s answer is slightly less accurate than Marit’s answer. ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... (1) (d) John then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane. C6H10(l) + H2(g) → C6H12(l) (i) Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene. ........................................................................................................................... ........................................................................................................................... (1) (ii) The percentage difference between these two methods (average bond enthalpies and enthalpies of combustion) is greater for cyclohexene than it was for ethene. John’s hypothesis was that it would be the same. Determine why the use of average bond enthalpies is less accurate for the cyclohexene equation shown above, than it was for ethene. Deduce what extra information is needed to provide a more accurate answer. ........................................................................................................................... ........................................................................................................................... ........................................................................................................................... 513 Version 2 2014 ........................................................................................................................... (2) (Total 7 marks) 14. Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of hydrogenation of ethene to form ethane. C2H4(g) + H2(g) → C2H6(g) John used the average bond enthalpies from Table 10. Marit used the values of enthalpies of combustion from Table 12. (a) Calculate the value for the enthalpy of hydrogenation of ethene obtained using the average bond enthalpies given in Table 10. ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... (2) (b) Determine the value for the enthalpy of hydrogenation of ethene using the values for the enthalpies of combustion of ethene, hydrogen and ethane given in Table 12. ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... (2) (c) Suggest one reason why John’s answer is slightly less accurate than Marit’s answer and calculate the percentage difference. ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... ...................................................................................................................................... (2) (d) John then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane. C6H10(l) + H2(g) → C6H12(l) (i) Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene. ........................................................................................................................... ........................................................................................................................... (1) 514 Version 2 2014 (ii) The percentage difference between these two methods (average bond enthalpies and enthalpies of combustion) is greater for cyclohexene than it was for ethene. John’s hypothesis was that it would be the same. Determine why the use of average bond enthalpies is less accurate for the cyclohexene equation shown above, than it was for ethene. Deduce what extra information is needed to provide a more accurate answer. ........................................................................................................................... ........................................................................................................................... ........................................................................................................................... ........................................................................................................................... (2) (Total 9 marks) 15. The standard enthalpy change of three combustion reactions is given below in kJ. 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(l) 2H2(g) + O2(g) → 2H2O(l) C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l) ∆HO = –3120 ∆HO = –572 ΔHO = –1411 Based on the above information, calculate the standard change in enthalpy, ∆HO, for the following reaction. C2H6(g) → C2H4(g) + H2 ............................................................................................................................................... . ............................................................................................................................................... . ............................................................................................................................................... . ............................................................................................................................................... . ............................................................................................................................................... . ............................................................................................................................................... . ............................................................................................................................................... . ............................................................................................................................................... . (Total 4 marks) 16. In some countries, ethanol is mixed with gasoline (petrol) to produce a fuel for cars called gasohol. (i) Define the term average bond enthalpy. (2) (ii) Use the information from Table 10 of the Data Booklet to determine the standard enthalpy change for the complete combustion of ethanol. 515 Version 2 2014 CH3CH2OH(g) + 3O2(g) → 2CO2(g) + 3H2O(g) (3) (iii) The standard enthalpy change for the complete combustion of octane, C8H18, is –5471 kJ mol–1. Calculate the amount of energy produced in kJ when 1 g of ethanol and 1 g of octane is burned completely in air. (2) (iv) Ethanol can be oxidized using acidified potassium dichromate, K2Cr2O7, to form two different organic products. CH3CH2OH A B State the structural formulas of the organic products A and B and describe the conditions required to obtain a high yield of each of them. (4) (v) Deduce and explain whether ethanol or A has the higher boiling point. (2) 516 Version 2 2014 (vi) Ethene can be converted into ethanol by direct hydration in the presence of a catalyst according to the following equation. C2H4(g) + H2O(g) CH3CH2OH(g) For this reaction identify the catalyst used and state one use of the ethanol formed other than as a fuel. (2) (Total 15 marks) Topic 5: Exam Questions Markscheme 1. D [1] 2. B [1] 3. D 4. D 5. C 6. C 7. C 8. A 9. D 10. B 11. A 12. D [1] [1] [1] [1] [1] [1] [1] [1] [1] [1] 13. (a) energy required = C=C + H–H/612 + 436 and energy released = C–C + 2(C–H)/347 +2(413) / energy required = C=C + H–H + 4(C–H)/612 + 436 + 4(413) and energy released = C–C + 6(C–H)/347 + 6(413); ∆H = (1048 – 1173)/(2700 – 2825) = –125 kJ mol–1; 2 (b) ∆H = –1411 + (–286) – (–1560) = –137 kJ mol–1; 1 (c) the actual values for the specific bonds may be different to the average 517 Version 2 2014 values / the combustion values referred to the specific compounds / OWTTE; (d) (i) –125 kJ mol–1; (ii) average bond enthalpies do not apply to the liquid state / OWTTE; the enthalpy of vaporization/condensation of cyclohexene and cyclohexane / OWTTE; 1 1 2 [7] 14. (a) (b) (c) (d) energy required = C=C + H–H/612 + 436 and energy released = C–C + 2(C–H)/347 + 2(413) / energy required = C=C + H–H + 4(C–H)/612 + 436 + 4(413) and energy released = C–C + 6(C–H)/347 + 6(413); ∆H = (1048 – 1173)/(2700 – 2825) = –125 kJ mol–1 2 ∆H = –1411 + (–286) – (–1560)/correct energy cycle drawn; = –137 kJ mol–1; Award [1 max] for incorrect or missing sign. 2 the actual values for the specific bonds may be different to the average values / the combustion values referred to the specific compounds / OWTTE; (137 125) (percentage difference) = × 100 = 8.76 %; 137 (137 125) Accept × 100 = 9.60 %. 125 (i) –125 kJ mol–1; (ii) average bond enthalpies do not apply to the liquid state / OWTTE; the enthalpy of vaporization/condensation of cyclohexene and cyclohexane / OWTTE; 2 1 2 [9] 15. (C2H6(g) + 3 12 O2(g) → 2CO2(g) + 3H2O(l)) ∆HO = –1560; (H2O(l) → H2(g) + ∆HO = +286; 1 2 O2(g)) (2CO2(g) + 2H2O(l) → C2H4(g) + 3O2(g)) (C2H6(g) → C2H4(g) + H2(g)) Allow other correct methods. Award [2] for –137. Allow ECF for the final marking point. ∆HO = +1411; ∆HO = +137(kJ); 4 [4] 16. (i) energy required to break (1 mol of) a bond in a gaseous molecule/state; Accept energy released when (1 mol of) a bond is formed in a gaseous molecule/state / enthalpy change when (1 mol of) bonds are made or broken in the gaseous molecule/state. average values obtained from a number of similar bonds/ compounds / OWTTE; 518 2 Version 2 2014 (ii) (iii) (iv) (v) (vi) Bonds broken (1)(C–C) + (1)(O–H) + (5)(C–H) + (1)(C–O) + (3)(O=O) = (1)(347) + (1)(464) + (5)(413) + (1)(358) + (3)(498) = 4728(kJ); Bonds formed (2 × 2)(C=O) + (3 × 2)(O–H) = (4)(746) + (6)(464) = 5768 (kJ); ∆H = 4728 – 5768 = –1040 kJ mol–1 / –1040 kJ; Units needed for last mark. Award [3] for final correct answer. Award [2] for +1040 kJ. 3 Mr(C2H5OH) = 46.08 / 46.1 and Mr(C8H18) = 114.26/114.3; 1 g ethanol produces 22.57 kJ and 1 g octane produces 47.88 kJ; Accept values ranges of 22.5–23 and 47.8–48 kJ respectively. No penalty for use of Mr = 46 and Mr = 114. 2 A: CH3CHO; B: CH3COOH/CH3CO2H; Accept either full or condensed structural formulas but not the names or molecular formulas. A: distillation; B: reflux; 4 ethanol/CH3CH2OH; hydrogen bonding (in ethanol); Award second point only if the first is obtained. 2 (concentrated) H3PO4 /(concentrated) phosphoric acid / H2SO4/sulfuric acid; dyes / drugs / cosmetics / solvent / (used to make) esters / (used in) esterification/disinfectant; 2 [16] Topic 15.1 Standard Enthalpy Changes of Reaction HL The enthalpy of a chemical reaction can be determine from standard enthalpy of formation tables The standard enthalpy of formation is the energy change that would result from the formation of a compound from its elements. These values have been compiled for a wide variety of substances and are tabulated. Since enthalpy values may vary with temperature and pressure of a gas, the standard tables are calibrated at 25 oC (298K) and 1 atmosphere. Enthalpy values also depend on the physical state of a compound. Since enthalpies of formation are defined as the difference in enthalpies of a compound and its elements, the standard enthalpies of formations for elements are assumed to be zero. The enthalpy change for a chemical reaction can be computed from this equation Hreaction = H products - H reactants Example Use the standard enthalpy of formation table to determine the enthalpy of reaction for this reaction: Ca(OH)2 (s) + 2 HCl (g ) CaCl2 (s) + 2 H2O (g) Enthalpies of formation: Ca(OH)2 (s) = - 986.2 kJ mol-1; 241.8 kJ mol-1; HCl (g ) = -92.30 kJ mol-1; CaCl2 (s) = -795.8 kJ mol-1; H2O (g) - Hreaction = [ 2(-241.8) + (-795.8)] -- [ (-986.2) + 2(-92.3)],kJ = (-483.6 -795.8) -(986.2-184.6) = (-1279.4) -(-1170.8) = -108.6 kJ 519 Version 2 2014 1. Use the standard enthalpy table to find the enthalpy of reaction for the following fuels A. Ethanol C2H5OH (g) + 3 O2 (g) 2 CO2 (g) + 3 H2O(g) B. Propane C. Benzene D. 2 C3H8 (g) + C6H6 (l) 5 O2 (g) + 7.5 O2 (g) 3 CO2 (g) + 6 CO2 (g) + 4 H2O (g) 3 H2O (g) On a per gram basis which of the above fuels produces the most heat? Calcium carbonate decomposes on heating as follows CaCO3 (s) CaO (s) + CO2 (g) Calculate the enthalpy of for this reaction 3 Hydrogen peroxide decomposes to water and oxygen according to the following reaction 2 H2O2 (l) 2 H2O (l) + O2 (g) Use this information and the enthalpies of formation from the table to calculate the enthalpy of this reaction. 4 Calcium carbide, CaC2, reacts with water to form acetylene C2H2, and Ca(OH)2 CaC2 (s) + 2 H2O (l) Ca(OH)2 (s) + C2H2 (g) ∆Hreaction = -127.2 kJ Use this information and the enthalpies of formation from the table to calculate the enthalpy of formation or calcium carbide. Enthalpy of Formation Calculate the enthalpy change for each of the following reactions. Label the reaction as exothermic or endothermic. 1) NH4NO3(s) NH4 + (aq) + NO3- (aq) –132 kJ 2+ 2) CaCl2(s) Ca (aq) –542.8 kJ 28.1 kJ Endothermic –205 kJ + 2 Cl (aq) –82.2 kJ Exothermic –167.2 kJ 3) 2 H2(g) + O2(g) 2 H2O(l) –571.6 kJ Exothermic 4) n-C4H10(g) + 13/2 O2(g) 4CO2(g) + 5H2O(l) -2878 kJExothermic Enthalpy of Formation 1) Calculate the enthalpy change for each of the following reactions. Label the reaction as exothermic or endothermic. a) C6H12O6(s) 2C2H5OH(l) + 2CO2(g) –74.4 kJ Exothermic b) n-C4H10(g) + 13/2 O2(g) 4CO2(g) + 5H2O(l) –2878 kJ Exothermic 2) Calculate the ΔHº for the reaction as well as when a 28 g sample of glucose, C 6H12O6(s), is burned to form CO2(g) and H2O(l) in a reaction at constant pressure and std pressure. ΔHrxn = –2808 kJ –434.4 kJ Exothermic 3) Calculate the ΔHº for the reaction as well as the amount of heat evolved or absorbed when a 0.1045 g sample of benzene, C6H6(l), is burned in excess oxygen to form CO2(g) and H2O(l) in a reaction at constant pressure and std. pressure. ΔHrxn = –3267.58 kJ –4.38 kJ Exothermic 520 Version 2 2014 521 Version 2 2014