CHAPTER 5 IB CHEMISTRY WARM UPS Mrs Hilliard
CHAPTER 5 IB CHEMISTRY WARM UPS Mrs. Hilliard
CHAPTER 5 VOCABULARY 1. Heat 10. Specific heat capacity 2. Entropy 11. Standard enthalpy change of formation 3. Enthalpy 4. System 5. Surroundings 6. Endothermic 7. Exothermic 8. Thermochemistry 9. Intensive property 12. Standard enthalpy change of combustion 13. Thermochemical equation 14. Bond enthalpy
1. 2. 3. 4. 5. 6. CHAPTER 5 VOCABULARY Heat- q, a form of energy that is transferred from a warmer 9. body to a cooler body as a result of the temperature gradient. 10. Entropy- a measure of the distribution of total available energy between particles. 11. Enthalpy- the heat content of a system at a constant pressure. Intensive property- a physical property that remains the same no matter how much of the substance is present. Specific heat capacity- the amount of heat required to raise 1 g of a substance by 1°C. Standard enthalpy change of formation- the energy change upon the formation of 1 mol of a substance from its constituent elements in their standard state. System- the part of the universe that contains the reaction 12. Standard enthalpy change of combustion- the heat or process being studied. evolved upon the complete combustion (burning) of 1 mol of substance. Surroundings- everything in the universe except the system. 13. Thermochemical equation- a balanced chemical equation that includes the physical states of all the reactants and the Endothermic- bond breaking, energy required to break the energy change, usually expressed as the change in bonds. enthalpy. 7. Exothermic- bond making, energy is released when new chemical bonds are made. 8. Thermochemistry- the study of heat changes that occur during chemical reactions. 14. Bond enthalpy- the energy required to break 1 mol of bonds in gaseous covalent molecules under standard conditions.
TYPES OF REACTIONS 1. Which statement about bonding is correct? 1. 2. 3. 4. Bond making is endothermic and releases energy Bond making is exothermic and requires energy Bond breaking is endothermic and releases energy 2. Some water is heated using the heat produced by the combustion of copper metal. Which values are need to calculate the enthalpy change of reaction? 1. The mass of copper 2. The change in temperature of the water 3. The mass of the water
TYPES OF REACTIONS 1. Which statement about bonding is correct? 1. 2. 3. 4. Bond making is endothermic and releases energy Bond making is exothermic and requires energy Bond breaking is endothermic and releases energy 2. Some water is heated using the heat produced by the combustion of copper metal. Which values are need to calculate the enthalpy change of reaction? 1. The mass of copper 2. The change in temperature of the water 3. The mass of the water
HEAT 3. The specific heat of lead is 0. 129 J g-1 K-1. What is the energy, in J, needed to increase the temperature of 100. 0 g of lead by 10. 0 K? 4. A pure calcium block with a mass of 20 g is heated so that its temperature increases from 10°C to 100°C. The specific heat capacity of calcium is 6. 74 x 10 -1 J g-1 K-1. Which expression gives the heat energy change in k. J? A. B. C. D. 20 x 6. 74 x 10 -1 J g-1 K-1 x 90 1000 20 x 6. 74 x 10 -1 J g-1 K-1 x 90 20 x 6. 74 x 10 -1 J g-1 K-1 x 363 1000 5. Which types of reactions are always exothermic? 1. Decomposition 2. Combustion 3. Neutralization
HEAT 3. The specific heat of lead is 0. 129 J g-1 K-1. What is the energy, in J, needed to increase the temperature of 100. 0 g of lead by 10. 0 K? Q=mc∆T Q= (100 g)(0. 129 J g-1 K-1)(10 K)=129 J 4. A pure calcium block with a mass of 20 g is heated so that its temperature increases from 10°C to 100°C. The specific heat capacity of calcium is 6. 74 x 10 -1 J g-1 K-1. Which expression gives the heat energy change in k. J? A. B. C. D. 20 x 6. 74 x 10 -1 J g-1 K-1 x 90 1000 20 x 6. 74 x 10 -1 J g-1 K-1 x 90 20 x 6. 74 x 10 -1 J g-1 K-1 x 363 1000 5. Which types of reactions are always exothermic? 1. Decomposition 2. Combustion
ENDOTHERMIC AND EXOTHERMIC REACTIONS 6. Which is correct about energy changes during bond breaking and bond formation? Bond breaking Bond formation Endothermic and ∆H positive Exothermic and ∆H negative Endothermic and ∆H negative Exothermic and ∆H positive Endothermic and ∆H negative Exothermic and ∆H negative Endothermic and ∆H positive 7. Which statement is correct given the enthalpy level diagram below? A. B. C. D. The reaction is exothermic and the reactants are more thermodynamically stable than the products. The reaction is endothermic and the products are more thermodynamically stable than the reactants. The reaction is endothermic and the reactants are more thermodynamically stable than the products. The reaction is exothermic and the products are more thermodynamically stable than the reactants.
ENDOTHERMIC AND EXOTHERMIC REACTIONS 6. Which is correct about energy changes during bond breaking and bond formation? Bond breaking Bond formation Endothermic and ∆H positive Exothermic and ∆H negative Endothermic and ∆H negative Exothermic and ∆H positive Endothermic and ∆H negative Exothermic and ∆H negative Endothermic and ∆H positive 7. Which statement is correct given the enthalpy level diagram below? A. B. C. D. The reaction is exothermic and the reactants are more thermodynamically stable than the products. The reaction is endothermic and the products are more thermodynamically stable than the reactants. The reaction is endothermic and the reactants are more thermodynamically stable than the products. The reaction is exothermic and the products are more thermodynamically stable than the reactants.
ENDOTHERMIC AND EXOTHERMIC REACTIONS 8. Which processes are exothermic? 1. Neutralization 2. Combustion 3. Ice melting 9. Which statement is correct given the enthalpy level diagram below? A. The reaction is exothermic and the reactants are more thermodynamically stable than the products. B. The reaction is endothermic and the products are more thermodynamically stable than the reactants. C. The reaction is endothermic and the reactants are more thermodynamically stable than the products. D. The reaction is exothermic and the products are more thermodynamically stable than the reactants.
ENDOTHERMIC AND EXOTHERMIC REACTIONS 8. Which processes are exothermic? 1. Neutralization 2. Combustion 3. Ice melting 9. Which statement is correct given the enthalpy level diagram below? A. The reaction is exothermic and the reactants are more thermodynamically stable than the products. B. The reaction is endothermic and the products are more thermodynamically stable than the reactants. C. The reaction is endothermic and the reactants are more thermodynamically stable than the products. D. The reaction is exothermic and the products are more thermodynamically stable than the reactants.
ENTHALPY Reaction ∆H A. Exothermic Positive B. Exothermic 10. Which is true for a chemical reaction in which the products have a higher C. enthalpy than the reactants? Negative Positive Endothermic 11. When some aqueous sodium hydroxide and aqueous hydrochloric acid D. were reacted together, the temperature of the surroundings was Endothermic observed to increase from 25°C to 30°C. What can be deduced from this observation? A. B. C. D. The reaction is endothermic and ∆H is negative. The reaction is endothermic and ∆H is positive. The reaction is exothermic and ∆H is negative. The reaction is exothermic and ∆H is positive. 12. When some solid ammonium thiosulfate and solid barium hydroxide were reacted together, the temperature of the surroundings was observed to decrease from 35°C to 21°C. What can be deduced from this observation? A. B. C. D. The reaction is endothermic and ∆H is negative. The reaction is endothermic and ∆H is positive. The reaction is exothermic and ∆H is negative. The reaction is exothermic and ∆H is positive. Negative
ENTHALPY Reaction ∆H A. Exothermic Positive B. Exothermic 10. Which is true for a chemical reaction in which the products have a higher C. enthalpy than the reactants? Negative Positive Endothermic 11. When some aqueous sodium hydroxide and aqueous hydrochloric acid D. were reacted together, the temperature of the surroundings was Endothermic observed to increase from 25°C to 30°C. What can be deduced from this observation? A. B. C. D. The reaction is endothermic and ∆H is negative. The reaction is endothermic and ∆H is positive. The reaction is exothermic and ∆H is negative. The reaction is exothermic and ∆H is positive. 12. When some solid ammonium thiosulfate and solid barium hydroxide were reacted together, the temperature of the surroundings was observed to decrease from 35°C to 21°C. What can be deduced from this observation? A. B. C. D. The reaction is endothermic and ∆H is negative. The reaction is endothermic and ∆H is positive. The reaction is exothermic and ∆H is negative. The reaction is exothermic and ∆H is positive. Negative
ENTHALPY 0. 5 g Al 13. Identical pieces of aluminum are added to two beakers, A and B, containing hydrochloric acid. Both acids have the same initial temperature but their volumes and concentrations differ. Which statement is correct? A. B. C. D. The temperature in A and B will increase at the same rate. It is not possible to predict whether A or B will have the higher maximum temperature. The maximum temperature in A will be higher than in B. The maximum temperature in A and B will be equal. 14. What is the energy, in k. J, released when 4. 00 mol of carbon monoxide is burned according to the following equation? 2 CO (g) + O 2 (g) → 2 CO 2 (g) ∆Hø = -564 k. J 15. How much energy, in joules, is required to increase the temperature of 4. 0 g of iron from 30 -38°C? (Specific heat of iron is 0. 450 J g-1 K-1).
ENTHALPY 0. 5 g Al 13. Identical pieces of aluminum are added to two beakers, A and B, containing hydrochloric acid. Both acids have the same initial temperature but their volumes and concentrations differ. Which statement is correct? A. B. C. D. The temperature in A and B will increase at the same rate. It is not possible to predict whether A or B will have the higher maximum temperature. The maximum temperature in A will be higher than in B. The maximum temperature in A and B will be equal. 14. What is the energy, in k. J, released when 4. 00 mol of carbon monoxide is burned according to the following equation? 2 CO (g) + O 2 (g) → 2 CO 2 (g) ∆Hø = 564 k. J 2 x 564= 1128 15. How much energy, in joules, is required to increase the temperature of 4. 0 g of iron from 30 -38°C? (Specific heat of iron is 0. 450 J g-1 K-1). Q=mc∆T Q= (4. 0 g)(0. 450 J g-1 K-1)(8°C)=14. 4 J
HESS’S LAW 16. Consider the following reactions. Cu 2 O (s) + ½ O 2 (g) → 2 Cu. O (s) ∆Hø= -163 k. J Cu 2 O (s) → Cu (s) + Cu. O (s) ∆Hø= +21 k. J What is the value of ∆Hø, in k. J, for this reaction? Cu (s) + ½ O 2 (g) → Cu. O (s) 17. Consider the following reactions. 2 Fe (s) + O 2 (g) → 2 Fe. O (s) ∆Hø= -521 k. J 4 Fe (s) + 3 O 2 (g) → 2 Fe 2 O 3 (s) ∆Hø= -1561 k. J What is the enthalpy change, in k. J, for the reaction below? 4 Fe. O (s) + O 2 (g) → 2 Fe 2 O 3 (s) 18. Which processes have a negative enthalpy change? A. H 2 O (g) → H 2 O (l) B. HCl (aq) + Na. OH (aq) → Na. Cl (aq) + H 2 O (l) C. 2 CH 3 OH (l) + 3 O 2 (g) → 2 CO 2 (g) + 4 H 2 O (l)
HESS’S LAW 16. Consider the following reactions. Cu 2 O (s) + ½ O 2 (g) → 2 Cu. O (s) ∆Hø= -163 k. J Cu 2 O (s) → Cu (s) + Cu. O (s) ∆Hø= +21 k. J What is the value of ∆Hø, in k. J, for this reaction? Cu (s) + ½ O 2 (g) → Cu. O (s) ∆Hø= -163 -21= -184 17. Consider the following reactions. 2 Fe (s) + O 2 (g) → 2 Fe. O (s) ∆Hø= -521 k. J 4 Fe (s) + 3 O 2 (g) → 2 Fe 2 O 3 (s) ∆Hø= -1561 k. J What is the enthalpy change, in k. J, for the reaction below? 4 Fe. O (s) + O 2 (g) → 2 Fe 2 O 3 (s) ∆Hø= 1561 – 2(-521) 18. Which processes have a negative enthalpy change? (exotherm) A. H 2 O (g) → H 2 O (l) Energy given off (condensation) React. higher than products for energy B. HCl (aq) + Na. OH (aq) → Na. Cl (aq) + H 2 O (l) (neutralization)
ENTHALPY CHANGES 19. When 200 cm 3 of 2. 0 mol dm-3 HCl is mixed with 200 cm 3 of 2. 0 mol dm-3 KOH, the temperature of the resulting solution increases by 10°C. What will be the temperature change in °C, when 100 cm 3 of these two solutions are mixed? 20. Consider the following reactions. N 2 (g) + O 2 (g) → 2 NO (g) ∆Hø=+150 k. J 2 NO 2 (g) → 2 NO (g) + O 2 (g) ∆Hø= +120 k. J What is the ∆Hø value, in k. J, for the following reaction? N 2 (g) + 2 O 2 (g) → 2 NO 2 (g) 21. The standard enthalpy changes for the combustion of carbon and carbon monoxide are shown below. C (s) + O 2 (g) → CO 2 (g) ∆Hø=-350 k. J CO (g) + ½ O 2 (g) → CO 2 (g) ∆Hø=-250 k. J What is the standard enthalpy change, in k. J, for the following reaction? C (s) + ½ O 2 (g) → CO (g)
ENTHALPY CHANGES 19. When 200 cm 3 of 2. 0 mol dm-3 HCl is mixed with 200 cm 3 of 2. 0 mol dm-3 KOH, the temperature of the resulting solution increases by 10°C. What will be the temperature change in °C, when 100 cm 3 of these two solutions are mixed? 10°C, ratios are the same, same reactants and products so energy doesn’t change 20. Consider the following reactions. N 2 (g) + O 2 (g) → 2 NO (g) ∆Hø=+150 k. J 2 NO 2 (g) → 2 NO (g) + O 2 (g) ∆Hø= +120 k. J What is the ∆Hø value, in k. J, for the following reaction? N 2 (g) + 2 O 2 (g) → 2 NO 2 (g) ∆Hø= 1 x (+150) + -1(+120)= 30 k. J 21. The standard enthalpy changes for the combustion of carbon and carbon monoxide are shown below. C (s) + O 2 (g) → CO 2 (g) ∆Hø=-350 k. J CO (g) + ½ O 2 (g) → CO 2 (g) ∆Hø=-250 k. J What is the standard enthalpy change, in k. J, for the following reaction? C (s) + ½ O 2 (g) → CO (g) ∆Hø=-350+ 250 = -100 k. J mol -1
BOND ENTHALPY 22. Using the equation below: C (s) + O 2 (g) → CO 2 (g) ∆Hø=-400 k. J mol -1 Mn (s) + O 2 (g) → Mn. O 2 (s) ∆Hø= -520 k. J mol -1 What is ∆H, in k. J, for the following reaction? Mn. O 2 (s) + C (s) → Mn (s) + CO 2 (g) 23. 2. 0 g of sodium hydroxide, Na. OH, was added to 98. 0 g of water. The temperature of the solution increased from 12. 0 °C to 17. 5°C. The specific heat capacity of the solution is 4. 18 J g-1 K-1. Give the expression for the heat evolved in k. J mol-1. 24. Which equation represent the bond enthalpy for the H- Cl bond in hydrogen chloride? A. B. C. D. HCl (g)→ H (g) + Cl (l) HCl (g)→ H (g) + Cl (g) HCl (g)→ H (g) + ½ Cl 2 (g) HCl (g)→ H (g) + ½ Cl (l)
BOND ENTHALPY 22. Using the equation below: C (s) + O 2 (g) → CO 2 (g) ∆Hø=-400 k. J mol -1 Mn (s) + O 2 (g) → Mn. O 2 (s) ∆Hø= -520 k. J mol -1 What is ∆H, in k. J, for the following reaction? Mn. O 2 (s) + C (s) → Mn (s) + CO 2 (g) ∆Hø= -400 + 520= 120 k. J 23. 2. 0 g of sodium hydroxide, Na. OH, was added to 98. 0 g of water. The temperature of the solution increased from 12. 0 °C to 17. 5°C. The specific heat capacity of the solution is 4. 18 J g-1 K-1. Give the expression for the heat evolved in k. J mol-1. 5. 5°C x 100. 0 g x 4. 18 J g-1 °C-1 x 40 g mol-1 / 1000 (conversion from J to k. J) 24. Which equation represent the bond enthalpy for the H- Cl bond in hydrogen chloride? A. B. C. D. HCl (g)→ H (g) + Cl (l) HCl (g)→ H (g) + Cl (g) HCl (g)→ H (g) + ½ Cl 2 (l)
RXN RATE 25. Consider the reaction between zinc and hydrochloric acid. Which factors will affect the reaction rate? A. The number of reactant particles that collide with the appropriate geometry B. The collision frequency of the reactant particles C. The number of reactant particles with E > Ea 26. In a reaction that occurs in 100 g of aqueous solution, the temperature of the reaction mixture increases by 15°C. If 0. 25 mol of the limiting reagent is consumed, what is the enthalpy change (in k. J mol -1) for the reaction? Assume the specific heat capacity of the solution = 3. 7 k. J kg -1 K -1. 27. Which equation best represents the bond enthalpy of HCl? A. B. C. D. 2 HCl (g) → H 2 (g) + Cl 2 (g) HCl (g) → ½ H 2 (g) + ½ Cl 2 (g) HCl (g) → H+ (g) + Cl- (g) HCl (g) → H (g) + Cl (g)
RXN RATE 25. Consider the reaction between zinc and hydrochloric acid. Which factors will affect the reaction rate? A. The number of reactant particles that collide with the appropriate geometry B. The collision frequency of the reactant particles C. The number of reactant particles with E > Ea 26. In a reaction that occurs in 100 g of aqueous solution, the temperature of the reaction mixture increases by 15°C. If 0. 25 mol of the limiting reagent is consumed, what is the enthalpy change (in k. J mol -1) for the reaction? Assume the specific heat capacity of the solution = 3. 7 k. J kg -1 K -1. -0. 100 kg x 3. 7 k. J kg -1 K -1 x 15°C/ 0. 25 mol 27. Which equation best represents the bond enthalpy of HCl? A. 2 HCl (g) → H 2 (g) + Cl 2 (g) B. HCl (g) → ½ H 2 (g) + ½ Cl 2 (g) C. HCl (g) → H+ (g) + Cl- (g)
RXNS 28. Which combination is correct for a chemical reaction that absorbs heat from the surroundings? 29. Which combination is correct for a chemical reaction that releases heat into the surroundings? 30. Which process represents the C - Cl bond enthalpy in tetrachloromethane? A. CCl 4 (l) → C (g) + 4 Cl (g) B. CCl 4 (l) → C (g) + 2 Cl 2 (g) C. CCl 4 (g) → CCl 3 (g) + Cl (g) Type of reaction ∆H at constant pressure A. Endothermic Positive B. Endothermic Negative C. Exothermic Positive D. Exothermic Type of reaction Negative ∆H at constant pressure A. Endothermic Positive B. Endothermic Negative C. Exothermic Positive D. Exothermic Negative
RXNS 28. Which combination is correct for a chemical reaction that absorbs heat from the surroundings? 29. Which combination is correct for a chemical reaction that releases heat into the surroundings? 30. Which process represents the C - Cl bond enthalpy in tetrachloromethane? A. CCl 4 (l) → C (g) + 4 Cl (g) B. CCl 4 (l) → C (g) + 2 Cl 2 (g) C. CCl 4 (g) → CCl 3 (g) + Cl (g) Type of reaction ∆H at constant pressure A. Endothermic Positive B. Endothermic Negative C. Exothermic Positive D. Exothermic Type of reaction Negative ∆H at constant pressure A. Endothermic Positive B. Endothermic Negative C. Exothermic Positive D. Exothermic Negative
RXNS 31. Use the average bond enthalpies below to calculate the enthalpy change, in k. J, for the following reaction. Bond energy/ k. J mol-1 H 2 (g) + I 2 (g) → 2 HI (g) I-I 151 H-H 436 H-I 298 32. Use the average bond enthalpies below to calculate the enthalpy change, in k. J, for the following reaction. Bond energy/ k. J H 2 (g) + Br 2 (g) → 2 HBr (g) mol-1 H-H 436 Br- Br 193 H- Br 33. Which of the following reactions are endothermic? A. CH 4 + 2 O 2 → CO 2 + 2 H 2 O B. Na. OH + HCl → Na. Cl + H 2 O C. Cl 2 → 2 Cl 366
RXNS 31. Use the average bond enthalpies below to calculate the enthalpy change, in k. J, for the following reaction. Bond energy/ k. J mol-1 H 2 (g) + I 2 (g) → 2 HI (g) I-I - ∑(BE bonds 151 ∑(BE bonds broken) H-H – 2 (298) 436 formed) (436 + 151) = -9 H-I 298 32. Use the average bond enthalpies below to calculate the Bond energy/ k. J enthalpy change, in k. J, for the following reaction. mol-1 H 2 (g) + Br 2 (g) → 2 HBr (g) H-H 436 ∑(BE bonds broken) - ∑(BE Br- Brbonds formed) 193 (436 + 193) – 2 (366) H- Br= -103 366 33. Which of the following reactions are endothermic? A. CH 4 + 2 O 2 → CO 2 + 2 H 2 O
WRITTEN TEST 34. The data to the right is from an experiment used to measure the enthalpy change for the combustion of 1 mole of glucose, C 6 H 12 O 6 (s). The timetemperature data was taken from a datalogging software program. Mass of sample of glucose, m= 0. 3251 g Heat capacity of the system, Csystem= 2. 713 k. J K-1. Calculate ∆T, for the water, surrounding the chamber in the calorimeter. 35. Determine the amount, in moles, of glucose. 36. Calculate the enthalpy change for the combustion of 1 mole of glucose. 37. Using table 13 of the data booklet, calculate the percentage experimental error based on the data used in this experiment. 24. 5 24. 0 23. 5 23. 0
WRITTEN TEST 34. The data to the right is from an experiment used to measure the enthalpy change for the combustion of 1 mole of glucose, C 6 H 12 O 6 (s). The time-temperature data was taken from a data- logging software program. Mass of sample of glucose, m= 0. 3251 g Heat capacity of the system, Csystem= 2. 713 k. J K-1. Calculate ∆T, for the water, surrounding the chamber in the calorimeter. Final – initial = 24. 8°C- 23. 0°C= 1. 8°C 35. Determine the amount, in moles, of glucose. 0. 3251 g x 1 mol/ 180. 18 g= 0. 001804 mol or 1. 804 x 10 -3 mol 36. Calculate the enthalpy change for the combustion of 1 mole of glucose. ∆Hc= c∆T/mol ∆Hc= 2. 713 k. J K-1)(1. 8°C)/ 0. 001804 mol ∆Hc= -2706. 98 k. J/mol or -2. 7 x 10 -3 k. J/mol 37. Using table 13 of the data booklet, calculate the percentage experimental error based on the data used in this experiment. % error= (theoretical)- (experimental) / theoretical x 100 24. 5 24. 0 23. 5 23. 0
WRITTEN TEST 38. A hypothesis is suggested that TNT, 2 -methyl-1, 3, 5 - trinitrobenzene, is a powerful explosive because it has: A large enthalpy of combustion A large volume of gas generated upon combustion 24. 5 24. 0 23. 5 A high reaction rate Use your answer in 36 and the following data to evaluate this hypothesis. Equation for combustion Relative rate of combustion Glucose C 6 H 12 O 6 (s) + 6 O 2 (g) → 6 CO 2 (g) + 6 H 2 O (g) Low TNT 2 C 7 H 5 N 3 O 6 (s) → 7 C(s) + 7 CO (g) + 3 N 2 (g) + 5 H 2 O (g) High 23. 0 Enthalpy of combustion/ k. J mol-1 3406
WRITTEN TEST 38. A hypothesis is suggested that TNT, 2 -methyl-1, 3, 5 - trinitrobenzene, is a powerful explosive because it has: A large enthalpy of combustion A large volume of gas generated upon combustion A high reaction rate Use your answer in 36 and the following data to evaluate this hypothesis. 1. Enthalpy of combustion is higher in glucose than TNT so enthalpy is not an important part of the explosive power. 2. The amount of gas (CO 2 + H 2 O) in glucose is higher in mol than the amount of gas (CO 2 + H 2 O + N 2) in TNT 3. High relative. Equation rate of combustion for TNT is. Relative important than glucose. for combustion rateand of higher Enthalpy of combustion Glucose C 6 H 12 O 6 (s) + 6 O 2 (g) → 6 CO 2 (g) + 6 H 2 O (g) Low TNT 2 C 7 H 5 N 3 O 6 (s) → 7 C(s) + 7 CO (g) + 3 N 2 (g) + 5 H 2 O (g) High combustion/ k. J mol-1 3406
ENTHALPY OF METHANOL 39. In an experiment to measure the enthalpy change of combustion of methanol, a student heated a copper calorimeter containing 200 cm 3 of water with a spirit lamp and collected the following data. Initial temperature of water: Final temperature of water: Mass of methanol burned: Density of water: 12°C 22°C 1. 32 g 1. 00 g cm-3 Use the data to calculate the heat evolved when the methanol was combusted. 40. Calcu. Iate the enthalpy change of combustion per mole of methanol. 41. Suggest two reasons why the result is not the same as the value in the Data Booklet.
ENTHALPY OF METHANOL 39. In an experiment to measure the enthalpy change of combustion of methanol, a student heated a copper calorimeter containing 200 cm 3 of water with a spirit lamp and collected the following data. Initial temperature of water: Final temperature of water: Mass of methanol burned: Density of water: 12°C 22°C 1. 32 g 1. 00 g cm-3 Use the data to calculate the heat evolved when the methanol was combusted. Q=mc∆T Q= (200 g)(4. 18 J g-1 K-1)(10°C)= 8360 J 40. Calcu. Iate the enthalpy change of combustion per mole of methanol. 1. 32 g methanol/ 32. 05 g/mol CH 3 OH= 0. 0412 mol CH 3 OH ∆H= Q/mol ∆H= 8360 J/0. 0412 mol = 202912. 62 J/mol or 202. 91 k. J/mol 41. Suggest two reasons why the result is not the same as the value in the Data Booklet. 1. Heat loss 2. Incomplete combustion 3. Heat absorbed by calorimeter not included
DIAZENE AND ETHANOL 42. The equation for the reaction between diazene and oxygen is given below. 2 N 2 H 2 (g) + O 2 (g) → 2 N 2 (g) + 2 H 2 O (g) Use the bond enthalpy values from Table 11 of the Data Booklet to determine the enthalpy change for this reaction. 43. In some countries, ethanol is mixed with gasoline to produce a fuel for cars called gasohol. Define the term average bond enthalpy. 44. Use the information in Table 11 of the Data Booklet to determine the standard enthalpy change for the complete combustion of ethanol. CH 3 CH 2 OH (g) + 3 O 2 (g) → 2 CO 2 (g) + 3 H 2 O (g) 45. The standard enthalpy change for the complete combustion of propane, C 3 H 8, is -2219 k. J mol-1. Calculate the amount of energy produced in k. J when 1 g of ethanol and 1 g of propane is burned
DIAZENE AND ETHANOL 42. The equation for the reaction between diazene and oxygen is given below. 2 N 2 H 2 (g) + O 2 (g) → 2 N 2 (g) + 2 H 2 O (g) Use the bond enthalpy values from Table 11 of the Data Booklet to determine the enthalpy change for this reaction. 4 N-H = 4 x 391= 1564, 2 N=N = 2 x 470= 940, 1 O=O = 498. Total bonds broken 3002. Bonds formed 2 N=N = 2 x 470= 940, 4 O-H = 4 x 463 = 1852. Total bonds formed 2792. 3002 -2792= +210 k. J/mol 43. In some countries, ethanol is mixed with gasoline to produce a fuel for cars called gasohol. Define the term average bond enthalpy. The energy required to break 1 mol of bonds in gaseous covalent molecules. It is averaged for the same bond in a number of similar compounds. 44. Use the information in Table 11 of the Data Booklet to determine the standard enthalpy change for the complete combustion of ethanol. CH 3 CH 2 OH (g) + 3 O 2 (g) → 2 CO 2 (g) + 3 H 2 O (g) 5 C-H = 5 x 414= 2070, 1 O-H = 463, 1 C-C = 346, 1 C- O = 358, 3 O=O = 3 x 498= 1494. Total bonds broken 4731. Bonds formed 4 C=O = 4 x 804= 3216, 6 O -H = 6 x 463 = 2778. Total bonds formed 5994. 4731 -5994= -1263 k. J/mol 45. The standard enthalpy change for the complete combustion of propane, C 3 H 8, is -2219 k. J mol 1. Calculate the amount of energy produced in k. J when 1 g of ethanol and 1 g of propane is burned completely in air. -2219 k. J mol-1/ 44. 11 g mol-1 = -50. 31 k. J (since both are 1 g) -1367 k. J mol-1/ 46. 08 g mol-1 = -29. 67 k. J Both answers must be given
CUSO 4 46. The data below are from an experiment to measure the enthalpy change for the reaction of aqueous copper (II) sulfate, Cu. SO 4 (aq) and zinc, Zn (s). Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq) 100 cm 3 of 1. 00 mol dm 3 copper (II) sulfate solution was placed in a polystyrene cup and zinc powder was added after 120 seconds. The temperature- time data was taken from a data logging software program. The initial 20 readings were at 19. 5°C. A straight line has been drawn through some of the data points. The equation for this line is given by the data logging software as T= -0. 100 t + 26. 5 where T is the Temperature at time t. The heat produced by the reaction can be calculated from the temperature change, ∆T, using the expression below. Heat change= Volume of Cu. SO 4 (aq) x Specific heat capacity of H 2 O x ∆T Describe two assumptions made in using this expression to calculate heat changes. Linear fit for selected data. T= -0. 100 t + 26. 5 T Temperature t time
CUSO 4 46. The data below are from an experiment to measure the enthalpy change for the reaction of aqueous copper (II) sulfate, Cu. SO 4 (aq) and zinc, Zn (s). Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq) 100 cm 3 of 1. 00 mol dm-3 copper (II) sulfate solution was placed in a polystyrene cup and zinc powder was added after 120 seconds. The temperature- time data was taken from a data logging software program. The initial 20 readings were at 19. 5°C. A straight line has been drawn through some of the data points. The equation for this line is given by the data logging software as T= -0. 100 t + 26. 5 where T is the Temperature at time t. The heat produced by the reaction can be calculated from the temperature change, ∆T, using the expression below. Heat change= Volume of Cu. SO 4 (aq) x Specific heat capacity of H 2 O x ∆T Describe two assumptions made in using this expression to calculate heat changes. 1. No heat is lost 2. The specific heat capacity of zinc is zero or negligible 3. Linear fit for selected data. T= -0. 100 t + 26. 5 T Temperature t time
CUSO 4 47. Use the data presented by the data logging software to deduce the temperature change, ∆T, which would have occurred if the reaction had taken place instantaneously with no heat loss. 48. State the assumption made in #47. 49. Calculate the heat, in k. J, produced during the reaction using the expression given in #46. 50. The color of the solution changed from blue to colorless. Deduce the amount, in moles, of zinc which reacted in the polystyrene cup. 51. Calculate the enthalpy change, in k. J mol-1, for this reaction. Linear fit for selected data. T= -0. 100 t + 26. 5 T Temperature t time
CUSO 4 47. Use the data presented by the data logging software to deduce the temperature change, ∆T, which would have occurred if the reaction had taken place instantaneously with no heat loss. (final temp – initial temp) 26. 5 - 19. 5= 7°C 48. State the assumption made in #47. Temperature decreases at a uniform rate (when above room temperature) 49. Calculate the heat, in k. J, produced during the reaction using the expression given in #46. 100 g x 4. 18 J g-1 K-1 x 7°C= 2926 J= 2. 9 k. J 50. The color of the solution changed from blue to colorless. Deduce the amount, in moles, of zinc which reacted in the polystyrene cup. n=M x dm 3 n= 1. 00 mol dm-3 x 0. 100 dm 3 = 0. 100 mol 51. Calculate the enthalpy change, in k. J mol-1, for this reaction. 2. 9 k. J/ 0. 100 mol = 29 k. J/mol Linear fit for selected data. T= -0. 100 t + 26. 5 T Temperature t time
METHANOL 52. Methanol is made in large quantities as it is used in the production of polymers and in fuels. The enthalpy of combustion of methanol can be determined theoretically or experimentally. CH 3 OH (l) + 1½ O 2 (g) → CO 2 (g) + 2 H 2 O (g) Using the information from Table 11 of the Data Booklet, determine theoretical enthalpy of combustion of methanol. 53. The enthalpy of combustion of methanol can also be determined experimentally in a school laboratory. A burner containing methanol was weighed and used to heat water in a test tube as illustrated below. The following data were collected. Calculate the amount, ofand methanol Initial massinofmol, burner methanol/burned. 95. 371 g Final mass of burner and methanol/ g 90. 538 Mass of water in test tube/g 20. 000 Initial temperature of water/ °C 23. 4°C Final temperature of water/ °C 41. 8°C
METHANOL 52. Methanol is made in large quantities as it is used in the production of polymers and in fuels. The enthalpy of combustion of methanol can be determined theoretically or experimentally. CH 3 OH (l) + 1½ O 2 (g) → CO 2 (g) + 2 H 2 O (g) Using the information from Table 11 of the Data Booklet, determine theoretical enthalpy of combustion of methanol. 3 C-H = 3 x 414= 1242, 1 C-O = 358, 1 O-H = 463, 1. 5 O=O = 1. 5 x 498= 747. Total bonds broken 2810. Bonds formed 2 C=O = 2 x 804= 1608, 4 O-H = 4 x 463 = 1852. Total bonds formed 3460. 2810 -3460 = -650 k. J/mol 53. The enthalpy of combustion of methanol can also be determined experimentally in a school laboratory. A burner containing methanol was weighed and used to heat water in a test tube as illustrated below. The following data were collected. Calculate the Initial mass of burner and 95. 371 amount, in mol, of methanol burned. 4. 833 g methanol/ g methanol x 1 mol/ Final mass of burner and methanol/ g 90. 538 32. 05 g methanol = Mass of water in test tube/g 20. 000 0. 151 mol burned Initial temperature of water/ °C 23. 4°C
METHANOL 54. Calculate the heat absorbed, in k. J, by the water. 55. Determine the enthalpy change, in k. J mol-1, for the combustion of 1 mole of methanol. 56. The Data Booklet value for the enthalpy of combustion of methanol is -726 k. J mol-1. Suggest why this value differs from the values calculated in #52 and #55. Initial mass of burner and methanol/ 95. 371 g Final mass of burner and methanol/ g 90. 538 Mass of water in test tube/g 20. 000 Initial temperature of water/ °C 23. 4°C Final temperature of water/ °C 41. 8°C
METHANOL 54. Calculate the heat absorbed, in k. J, by the water. Q= mc∆T 20 g x 4. 18 J g-1 K-1 x 18. 4°C= 1538. 24 J= 1. 54 k. J 55. Determine the enthalpy change, in k. J mol-1, for the combustion of 1 mole of methanol. -1. 54 k. J/0. 151 mol= -10. 2 k. J/mol (neg. due to exotherm rxn) 56. The Data Booklet value for the enthalpy of combustion of methanol is -726 k. J mol-1. Suggest why this value differs from the values calculated in #52 and #55. #52 - Bond enthalpies are average values and differ slightly from one compound to another depending Initialto mass burner and on neighboring atoms. #55 - Heat is lost theofsurroundings or 95. 37 not methanol/ g 1 all heat transferred to the water. Final mass of burner and methanol/ 90. 53 g 8 Mass of water in test tube/g 20. 00 0
PROPENE 57. Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of hydrogenation of propene to form propane. C 3 H 6 (g) + H 2 (g) → C 3 H 8 (g) Karen used the average bond enthalpies from Table 11. Jessica used the values of enthalpies of combustion from Table 13. Calculate the value for the enthalpy of hydrogenation of propene obtained using the average bond enthalpies given in Table 11. 2 O 2 -1131 k. J mol-1 3 O 2 -1379 k. J mol-1 58. Jessica arranged the values she found in Table 13 into an energy cycle. Calculate the value for the enthalpy of hydrogenation of propene from the energy cycle. ∆Hø (hydrogenation) C 3 H 6 (g) + H 2 (g) C 3 H 8 (g) 5 O 2 9 -221 l-1 o 2 CO k. J m 2 (g) + H 2 O(l)
PROPENE 57. Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of hydrogenation of propene to form propane. C 3 H 6 (g) + H 2 (g) → C 3 H 8 (g) Karen used the average bond enthalpies from Table 11. Jessica used the values of enthalpies of combustion from Table 13. Calculate the value for the enthalpy of hydrogenation of propene obtained using the average bond enthalpies given in Table 11. 1 C-C= 346, 1 C=C = 614, 6 C-H = 6 x 414= 2484, 1 H-H = 436. Total bonds broken 3880. 2 C-C = 2 x 346= 692, 8 C-H = 8 x 414 = 3312. Total bonds formed 4004. 3880 -4004= -124 k. J/mol 2 O 2 -1131 k. J mol-1 3 O 2 -1379 k. J mol-1 58. Jessica arranged the values she found in Table 13 into an energy cycle. Calculate the value for the enthalpy of hydrogenation of propene from the energy cycle. ∆Hø (hydrogenation) C 3 H 6 (g) + H 2 (g) C 3 H 8 (g) 5 O 2 219 -1 -2+ ol k. J m 3 CO 2 (g) 4 H 2 O(l) (-1379+ -1131) – (-2219)= -291 k. J/ mol
PROPANE 59. Suggest one reason why Karen’s answer is slightly less accurate than Jessica’s answer. 60. Karen then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane. Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene. C 6 H 10 (l) + H 2 (g) → C 6 H 12 (l) 61. The percentage difference between these two methods (average bond enthalpies and enthalpies of combustion) is greater for cyclohexene than it was for propene. Karen’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 propene. Deduce what extra information is needed to provide a more accurate answer.
PROPANE 59. Suggest one reason why Karen’s answer is slightly less accurate than Jessica’s answer. The actual values for the specific bonds may be different from the average values. 60. Karen then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane. Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene. C 6 H 10 (l) + H 2 (g) → C 6 H 12 (l) 10 C-H= 10 x 414= 4140, 5 C-C = 5 x 346= 1730, 1 C=C = 507, 1 H-H =436 Total bonds broken= 6813 12 C-H= 12 x 414= 4968, 6 C-C = 6 x 346= 2076= 7044 (6813 -7044)= -231 k. J/mol 61. The percentage difference between these two methods (average bond enthalpies and enthalpies of combustion) is greater for cyclohexene than it was for propene. Karen’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 propene. Deduce what extra information is needed to provide a more accurate answer. 1. Average bond enthalpies do not apply to the liquid state. 2. The extra info needed is the enthalpy of
COPPER (I) SULFATE 62. If anhydrous copper (I) sulfate powder is left in the atmosphere it slowly absorbs water vapor giving the blue pentahydrated solid. Cu 2 SO 4 (s) + 5 H 2 O(l) → Cu 2 SO 4 ∙ 5 H 2 O(s) It is difficult to measure the enthalpy change for this reaction directly. However, it is possible to measure the heat changes directly when both anhydrous and pentahydrated copper (I) sulfate are separately dissolved in water, and then use an energy cycle to determine the required enthalpy change value, ∆Hx, indirectly. To determine ∆H 1 a student placed 15. 0 g of water in a cup made of expanded polystyrene and used a data logger to measure the temperature. After 2 minutes she dissolved 2. 37 g of anhydrous copper (I) sulfate in the water and continued to record the temperature while continuously stirring. She obtained the following results. Calculate the amount, in mol, of anhydrous copper (I) sulfate dissolved in the 15. 0 g of water. ∆Hx Cu 2 SO 4 (s) + 5 H 2 O(l) Cu 2 SO 4 ∙ 5 H 2 O(s) ∆H 1 ∆H 2 Cu 2 SO 4 (s)
COPPER (I) SULFATE 62. If anhydrous copper (I) sulfate powder is left in the atmosphere it slowly absorbs water vapor giving the blue pentahydrated solid. Cu 2 SO 4 (s) + 5 H 2 O(l) → Cu 2 SO 4 ∙ 5 H 2 O(s) It is difficult to measure the enthalpy change for this reaction directly. However, it is possible to measure the heat changes directly when both anhydrous and pentahydrated copper (I) sulfate are separately dissolved in water, and then use an energy cycle to determine the required enthalpy change value, ∆Hx, indirectly. To determine ∆H 1 a student placed 15. 0 g of water in a cup made of expanded polystyrene and used a data logger to measure the temperature. After 2 minutes she dissolved 2. 37 g of anhydrous copper (I) sulfate in the water and continued to record the temperature while continuously stirring. She obtained the following results. Calculate the amount, in mol, of anhydrous copper (I) sulfate dissolved in the 15. 0 g of water. ∆Hx Cu 2 SO 4 (s) + 5 H 2 O(l) Cu 2 SO 4 ∙ 5 H 2 O(s) ∆H 1 ∆H 2 Cu 2 SO 4 (s) 2. 37 g Cu 2 SO 4 x 1 mol/ 223. 11 g Cu 2 SO 4 = 0. 0106 mol Cu 2 SO 4
COPPER (I) SULFATE 63. Determine what the temperature rise would have been, in °C, if no heat had been lost to the surroundings. 64. Calculate the heat change, in k. J, when 2. 37 g of anhydrous copper (I) sulfate is dissolved in the water. 65. Determine the value of ∆H 1 in k. J mol -1. 66. To determine ∆H 2, 5. 34 g of pentahydrated copper (I) sulfate was dissolved in 13. 35 g of water. It was observed that the temperature of the solution decreased by 2. 30°C. Calculate the amount, in mol, of water in 5. 34 g of pentahydrated copper (I) sulfate.
COPPER (I) SULFATE 63. Determine what the temperature rise would have been, in °C, if no heat had been lost to the surroundings. Final temp- initial temp = 26. 1°C – 19. 1°C = 7 °C 64. Calculate the heat change, in k. J, when 2. 78 g of anhydrous copper (I) sulfate is dissolved in the water. Q= mc∆T Q= 15 g x 4. 18 J/g°C x 7°C= 438. 9 J = 0. 44 k. J or 17. 78 g x 4. 18 J/g°C x 7°C= 520. 2428 J = 0. 52 k. J 65. Determine the value of ∆H 1 in k. J mol -1. 0. 44 k. J/0. 0106 mol= -41. 5 k. J/mol (exothermic or temp. rose) or 0. 52 k. J/ 0. 0106 mol= -49. 1 k. J/mol 66. To determine ∆H 2, 5. 34 g of pentahydrated copper (I) sulfate was dissolved in 13. 35 g of water. It was observed that the temperature of the solution decreased by 2. 30°C. Calculate the amount, in mol, of water in 5. 34 g of pentahydrated copper (I) sulfate. 5. 34 g Cu 2 SO 4 ∙ 5 H 2 O(s) /313. 21 Cu 2 SO 4 ∙ 5 H 2 O(s)= 0. 017 mol Cu 2 SO 4 ∙ 5 H 2 O
COPPER (I) SULFATE 67. Determine the value of ∆H 2 in k. J mol -1. 68. Using the values obtained for ∆H 1 in #65 and #67, determine the value for ∆Hx in k. J mol -1. 69. The magnitude (the value without the + or – sign) found in a data book for ∆Hx is 78. 0 k. J mol -1. Calculate the percentage error obtained in this experiment. (If you did not obtain an answer for the experimental value of ∆Hx then use the value 70. 0 k. J mol -1, but this is not the true value. ) 70. The student recorded in her qualitative data that the anhydrous copper (I) sulfate she used was pale blue rather than completely white. Suggest a reason why it might have had this pale blue color and deduce how this would have affected the value she obtained for ∆Hx.
COPPER (I) SULFATE 67. Determine the value of ∆H 2 in k. J mol -1. Q= mc∆T Q= 13. 35 g x 4. 18 J/g°C x 2. 3°C= 128. 35 J = 0. 13 k. J 0. 13/0. 017= 7. 6 k. J/mol or Q= 16. 13 g x 4. 18 J/g°C x 2. 3°C= 155. 07 J = 0. 16 k. J/ 0. 017 mol= 9. 4 k. J/mol 68. Using the values obtained for ∆H 1 in #65 and #67, determine the value for ∆Hx in k. J mol -1. ∆Hx = ∆H 1 - ∆H 2 ∆Hx = -41. 5 - 7. 6= -49. 1 69. The magnitude (the value without the + or – sign) found in a data book for ∆Hx is 68. 0 k. J mol -1. Calculate the percentage error obtained in this experiment. (If you did not obtain an answer for the experimental value of ∆Hx then use the value 60. 0 k. J mol -1, but this is not the true value. ) % error = theoretical – experimental/ theoretical x 100= -68 - (-49. 1)/ -68 x 100 = 29. 26% or -68 -(-60)/ -68 x 100 = 11. 76% 70. The student recorded in her qualitative data that the anhydrous copper (I) sulfate she used was pale blue rather than completely white. Suggest a reason why it might have had this pale blue color and deduce how this would have affected the value she obtained for ∆Hx. The anhydrous copper (I) sulfate absorbed water from the air. The value would be less exothermic or less negative than expected as
GLUCOSE 71. The data below is from an experiment used to measure the enthalpy change for the combustion of 1 mole of glucose C 6 H 12 O 6 (s). The time- temperature data was taken from a datalogging software program. Mass of sample of glucose, m= 0. 3653 g Heat capacity of the system, Csystem = 2. 35 k. J K -1 Calculate ∆T, for the water, surrounding the
GLUCOSE 71. The data below is from an experiment used to measure the enthalpy change for the combustion of 1 mole of glucose C 6 H 12 O 6 (s). The timetemperature data was taken from a data- logging software program. Mass of sample of glucose, m= 0. 3653 g Heat capacity of the system, Csystem = 2. 35 k. J K-1 Calculate ∆T, for the water, surrounding the chamber in the calorimeter. Final- initial= 23. 78 -22. 01= 1. 77°C
GLUCOSE 72. Determine the amount, in moles, of glucose. 73. Calculate the enthalpy change for the combustion of 1 mole of glucose. 74. Using Table 13 of the Data Booklet, calculate the percentage experimental error based on the data used in this experiment. 75. A hypothesis is suggested that TNT, 2 -methyl-1, 3, 5 - trinitrobenzene, is a powerful explosive because it has: A. A large enthalpy of combustion B. A high reaction rate C. A large volume of gas generated upon combustion Equation for combustion Relative rate of hypothesis: Enthalpy of Use your answer in # 73 and the following data to evaluate this combustion Glucos e C 6 H 12 O 6 (s) + 6 O 2 (g) → 6 CO 2 (g) + 6 H 2 O (g) Low combustion/ k. J mol-1
GLUCOSE 72. Determine the amount, in moles, of glucose. 0. 3653 g C 6 H 12 O 6 (s) x 1 mol/ 180. 18 g C 6 H 12 O 6 = 0. 002027 mol 73. Calculate the enthalpy change for the combustion of 1 mole of glucose. ∆Hc= c∆T/mol 2. 35 k. J K-1 x 1. 77°C / 0. 002027= -2052. 05 k. J/mol (combustion reaction always exothermic) 74. Using Table 13 of the Data Booklet, calculate the percentage experimental error based on the data used in this experiment. -2803 -(-2052. 05)/-2803 x 100= 26. 79% 75. A hypothesis is suggested that TNT, 2 -methyl-1, 3, 5 - trinitrobenzene, is a powerful explosive because it has: A. A large enthalpy of combustion B. A high reaction rate C. A large volume of gas generated upon combustion Use your answer in # 73 and the following data to evaluate this hypothesis: 1. Enthalpy of combustion is higher in glucose than TNT so enthalpy is not an important part of explosive power. Equation for combustion Relative rate of Enthalpy of 2. High relative rate of combustion for TNT is important and higher than glucose. combustion/ k. J 3. The volume of gas (CO 2 (g) + 6 H 2 O) in glucose is higher in mol than volume of gas (CO mol 2 -1 + H 2 O + N 2) 4. of TNT. Glucos C H O (s) + 6 O (g) → 6 CO (g) + 6 H O (g) Low e 6 12 6 2 2 2
COMBUSTION REACTIONS 76. The standard enthalpy change of three combustion reactions is given below in k. J. 2 C 2 H 6 (g) + 7 O 2 (g) → 4 CO 2 (g) + 6 H 2 O (l) ∆Hø = -3115 2 H 2 (g) + O 2 (g) → 2 H 2 O (l) ∆Hø = -537 C 2 H 4 (g) + 3 O 2 (g) → 2 CO 2 (g) + 2 H 2 O (l) ∆Hø = -1431 Based on the above information, calculate the standard change in enthalpy, ∆Hø, for the following reaction. C 2 H 6 (g) → C 2 H 4 (g) + H 2 (g)
COMBUSTION REACTIONS 76. The standard enthalpy change of three combustion reactions is given below in k. J. 2 C 2 H 6 (g) + 7 O 2 (g) → 4 CO 2 (g) + 6 H 2 O (l) ∆Hø = -3115/2 C 2 H 6 (g) + 3 ½ O 2 (g) → 2 CO 2 (g) + 3 H 2 O (l) 2 H 2 (g) + O 2 (g) → 2 H 2 O (l) ∆Hø = +537/2 H 2 O (l) → H 2 (g) + ½ O 2 (g) C 2 H 4 (g) + 3 O 2 (g) → 2 CO 2 (g) + 2 H 2 O (l) ∆Hø = +1431 2 CO 2 (g) + 2 H 2 O (l) → C 2 H 4 (g) + 3 O 2 (g) Based on the above information, calculate the standard change in enthalpy, ∆Hø, for the following reaction. C 2 H 6 (g) → C 2 H 4 (g) + H 2 (g)
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