Download AP Chemistry 2015-‐‑2016 Name: Chapter 5: Thermodynamics

AP Chemistry 2015-­‐‑2016 Chapter 5: Thermodynamics Name: Date: Per: 5.1-­‐‑5.4 Thermodynamic Basics 1) Define and describe the following thermochemistry terms: Helpful Information: a) Energy Ability to do work or transfer hear w = -­‐‑P∆V b) Work 101 J = 1 atm ∙ L Energy that causes an object to move against a force PV = nRT (Ideal Gas Law) c) Heat R values listed on your cheat sheet Energy transferred from hotter to colder object STP – standard conditions d) Kinetic energy Energy of motion e) Potential energy Energy of position relative to another object f) Joule Unit of energy. 1 J = 1 Kg*m2/s2 g) System What we re studying; typically the atoms/molecules involved in a reaction h) Surroundings Everything that s not the system i) Internal energy Sum of all energies of a system j) Endothermic System absorbs heat k) Exothermic System produces/releases heat l) State function Property of system that is determined by specifying conditions or its state 2) Thoroughly explain this equation for electrostatic potential energy: Eel = kQ1Q2 including each letter or symbol. d E= potential energy – interaction between two charged particles (cation and anion). k = proportionality constant: 8.99 x 109 J*m/C2. d = distance between atoms. Q = magnitude of electrical charges 3) Describe the First Law of Thermodynamics. Energy is not created or destroyed, just transferred or changed. 4) The First Law of Thermodynamics can be quantitatively summarized by the equation: ∆E = q + w a) What does q represent? Heat added to or lost from a system b) What does w represent? Work done by or on the system c) Under what conditions will the quantities q and w be negative numbers? Discuss each separately. Q is negative when heat is lost, w is negative when work is done by system on surroundings 5) Calculate ∆E and determine if the process is endothermic or exothermic for the following cases: WATCH YOUR UNITS!! a) A system absorbs 85 kJ of heat from its surroundings while doing 29 kJ of work on the surroundings. 56 kJ b) q = 1.50 kJ and w = –657 J 0.843 kJ c) The system releases 57.5 kJ of heat while doing 13.5 kJ of work on the surroundings. -­‐‑71 kJ 6) Calculate ∆E for each of the following cases: a) q = + 51 kJ, w = -­‐‑ 15 kJ 36 KJ b) q = + 100. kJ, w = -­‐‑ 65 kJ 35 KJ c) q = -­‐‑ 65 kJ, w = -­‐‑ 20 kJ -­‐‑85 KJ d) In which of these cases does the system do work on the surroundings? All of them – all have -­‐‑w 7) Calculate ∆E for each of the following: a) q = -­‐‑ 47 kJ, w = + 88 kJ 41 kJ b) q = + 82 kJ, w = + 47 kJ 129 kJ c) q = + 47 kJ, w = 0 47 kJ d) In which of these cases do the surroundings do work on the system? A+B = + w 8) A system releases 125 kJ of heat while 104 kJ of work is done on the system. Calculate the change in internal energy (in kJ). -­‐‑21 kJ 9) A system undergoes a process consisting of the following two steps: Step 1: The system absorbs 73 J of heat while 35 J of work is done on it. Step 2: The system absorbs 35 J of heat while performing 72 J of work. Calculate the change in internal energy for the overall process (in J). 71J 10) The volume of an ideal gas is decreased from 5.0 L to 5.0 mL at constant pressure of 2.0 atm. Calculate the work associated with this process (in J). WATCH YOUR UNITS!!! 1008.99 J 11) The reaction of nitrogen with hydrogen to make ammonia is N2 (g) + 3H2 (g) 2HN3 (g) ΔH = -­‐‑92.2 kJ What is the value of ΔE (in kJ) if the reaction is carried out at a constant pressure of 40.0 atm and the volume change is -­‐‑1.12L? WATCH YOUR UNITS!!! -­‐‑87.6752 kJ 12) A strip of magnesium of mass 15 g is dropped into a beaker of dilute hydrochloric acid. What work is done on the surrounding atmosphere (1.00 atm pressure, 25°C) by the subsequent reaction? -­‐‑1525 J 13) Calculate the work done (in joules) by a chemical reaction if the volume increases from 3.2 L to 3.4 L against a constant external pressure of 3.6 atm. What is the sign of the energy change? -­‐‑72.72 J 14) When solutions containing silver ions and chloride ions are mixed, silver chloride precipitates: Ag+(aq) + Cl–(aq) → AgCl(s) ∆H = –65.5 kJ a) Calculate ∆H for the formation of 2.00 mol of AgCl. -­‐‑131 kJ b) Calculate ∆H for the formation of 2.50 g of AgCl -­‐‑1.14 kJ 15) You are given ∆H for a process that occurs at constant pressure. What additional information is needed to determine ∆E for the process? Change in volume 16) A gas is confined to a cylinder with a piston under constant atmospheric pressure (fig. 5.3). When the gas reacts, it releases 79 kJ of heat to its surroundings and does 18 kJ of P-­‐‑V work on its surroundings. What are the values for ∆H and ∆E for this process? E = -­‐‑97 kJ; H = -­‐‑79 kJ (see pg 164 equation 5.10 for explanation) 17) The thermochemical equation for the burning of one mole of benzene under standard conditions is C6H6(l) + 15/2 O2(g) → 6 CO2(g) + 3 H2O(l) ∆H comb = –3267.7 kJ a. Is this reaction exothermic or endothermic? How do you know? exothermic b. How much heat is released when a 5.00-­‐‑g sample of benzene is burned in excess oxygen under standard conditions? (m.w. C6H6 = 78.11 g/mol) -­‐‑209 kJ 18) Used in welding metals, the reaction of acetylene with oxygen: C2H2(g) + 5/2 O2(g) →H2O(g) + 2 CO2(g) ∆H=–1255.5 kJ How much PV work is done (in kilojoules) and what is the value of ∆E (in kilojoules) for the reaction of 6.50 g of acetylene at atmospheric pressure if the volume change is –2.80 L? W = 0.2828 kJ; E = -­‐‑312.7 kJ 19) Aluminum metal reacts with chlorine with a spectacular display of sparks: 2 Al(s) + 3 Cl2(g) →2 AlCl3(s) ∆H= –1408.4 kJ How much heat (in kilojoules) is released on reaction of 5.00 g of Al? -­‐‑1408.4 kJ 5.5: Calorimetry 1. Water has a specific heat capacity of 4.184 J/g·°C. This means it takes 4.184 J to heat 1.00 gram of water 1.00°C. a) How much energy will it take to heat 10.0 grams of water 1°C? ______________ 41.84 J b) How much energy is needed to heat 30.0 g H2O from 10.0 °C to 50.0 °C? ____________ 5020 J 2. Let’s try a standard calorimetry problem. A pot of water (2.5 Liters of water) initially at 25.0°C is heated to boiling (100.°C). How much energy (in J) is needed to heat the water? (The density of water is 1 g/mL.) 780000 J What would this amount of heat be in kJ? ___________ 780 kJ 3. What amount of heat is released when 175 g of water cools from 100.°C to room temperature, 20.0 °C? 58600 J 4. We don’t always have to warm up or cool down water. The specific heat capacity of copper metal is 0.39 J/g·°C. It is _____________ (easier/more difficult) to heat up copper than to heat up water. How much energy would it take to heat up a 5.20 g sample of copper from 20.0 °C to 100.°C? 162 J 5. If 300. J of heat energy were used to heat up a 5.00 gram sample of copper metal and a 5.00 gram sample of water both starting at 10.0°C, calculate the final temperature of each sample? Water: 24.3˚C; Copper: 164˚C Signs of ΔT and q: − q means heat is released. + q means heat is absorbed. ΔT is always final temperature – initial temperature. If something is getting hotter (10° → 30°) the ΔT is 30 – 10 = + 20°. (heat is absorbed) If something is getting cooler (75° → 25°) the ΔT is 25 – 75 = − 50°. (heat is released) 6. Suppose we mix 90.0 grams of hot water (90.0°C) with 10.0 grams of cold water (10.0°C). Let x = the final temperature. C = 4.184 J/g·°C a. Set up an expression for the energy released (q) by the hot water (Δqhot = mhotCΔThot) q = 90*4.184*∆T; q = 90*4.184*(TF-­‐‑90) b. Set up an expression for the energy absorbed (q) by the cold water (Δqcold = mcoldCΔTcold) q = 10*4.184*(TF-­‐‑10) c. Knowing that the heat released = − heat absorbed, combine the two expressions and solve for x. 82˚C 7. We don’t always have to use water. Let’s use some aluminum shot (pellets). 175 grams of hot aluminum (100.°C) is dropped into an insulated cup that contains 40.0 mL of ice cold water (0.0°C). Follow the example above to determine the final temperature, x. a. Set up an expression for the heat lost by the aluminum (C=0.900 J/g·°C) b. Set up an expression for the heat gained by the cold water. c. Put the two expressions together (don’t forget to change one of the signs) and solve for x. 48˚C 8. Somewhat Confusing Definitions: There are several terms used in this chapter that sound very similar. Use the data provided to calculate each of them to clarify the differences. I’ve added some “Notes” that I hope will help. 74.8 J of heat is required to raise the temperature of 18.69 g of silver from 10.0°C to 27.0°C. a. What is the heat capacity of the silver sample? (J/°C) Note: This is a useful value only for this specific sample of silver. 4.4 J/˚C b. What is the specific heat capacity of silver? (J/g·°C) Note: This is a useful value for any sample of silver that is heated or cooled. This is equivalent to the 4.184 J·g-­‐‑1·°C-­‐‑1 that we use for water. This value is also called the specific heat. 0.235 J/g*˚C 5.6-­‐‑5.7: Hess’s Law and Enthalpies of Formation Use standard enthalpies of formation (appendix C) to determine the change in enthalpy for each reaction a) NaOH(s) + HCl(g) -­‐‑-­‐‑-­‐‑-­‐‑> NaCl(s) + H2O(g) -­‐‑133.9 kJ b) 2 CO(g) + O2(g) -­‐‑-­‐‑-­‐‑> 2 CO2(g) -­‐‑566 kJ c) CH4(g) + 2 O2(g) -­‐‑-­‐‑-­‐‑> CO2(g) + 2 H2O(l) -­‐‑890.3 kJ d) 2 H2S(g) + 3 O2(g) -­‐‑-­‐‑-­‐‑> 2 H2O(l) + 2 SO2(g) -­‐‑1123.6 kJ e) 2 NO(g) + O2(g) -­‐‑-­‐‑-­‐‑> 2 NO2(g) -­‐‑113 kJ 1. Using standard enthalpies of formation, calculate the standard enthalpy change for the combustion of 1 mol benzene, C6H6(l), to CO2 (g) and H2O (l). -­‐‑3267.4 kJ 2. Using standard enthalpies of formation, calculate the enthalpy change for the combustion of 1 mol of ethanol: C2H5OH (l) + 3O2 (g) → 2CO2 (g) + 3H2O (l) -­‐‑1366.7 kJ 3. The standard enthalpy change for the reaction CaCO3(s) → CaO(s) + CO2(g) is 178.1 kJ. From the values for the standard enthalpy of formation of CaO(s) and CO2(g), calculate the standard enthalpy of formation of CaCO3(s). -­‐‑1207.1 kJ 4. Given the following standard enthalpy of reaction, use the standard enthalpies of formation to calculate the standard enthalpy of formation of CuO(s): CuO(s) + H2(g) → Cu(s) + H2O(l) ΔH° = -­‐‑129.7 kJ -­‐‑156.1 kJ 5. Using Hess’s Law, calculate the standard enthalpy of formation of gaseous diborane (B2H6) using the following thermochemical information: 4B(s) + 3O2(g) → 2B2O3(s) ΔH = -­‐‑2509.1 kJ 2H2(g) + O2(g) → 2H2O(l) ΔH = -­‐‑571.7 kJ B2H6(g) + 3O2(g) → B2O3(s) + 3H2O(l) ΔH = -­‐‑2147.5 kJ 35.4 kJ 6. Naphthalene (C10H8) is a solid aromatic compound often sold as mothballs. The complete combustion of this substance to yield CO2(g) and H2O(l) at 25° C yields 5154 kJ/mol. a) Write balanced equations for the formation of naphthalene from the elements and for its combustion. b) Calculate the standard enthalpy of formation of naphthalene. C10H8 (s) + 12O2 (g) → 10CO2 (g) + 4H2O (l) 72.8 kJ 7. Iron ore can be converted to iron metal with CO gas. FeO (s) + CO (g) → Fe (s) + CO2 (g)
Calculate the standard enthalpy change for this reaction from these reactions
of iron oxides with CO :
(1) 3 Fe2O3 (s) + CO (g) → 2 Fe3O4 (s) + CO2 (g)
ΔH° = - 47 kJ
(2) Fe2O3 (s) + 3 CO (g) → 2 Fe (s) + 3 CO2 (g)
(3) Fe3O4 (s) + CO (g) → 3 FeO (s) + CO2 (g)
ΔH° = - 25 kJ
ΔH° = 19 kJ
-­‐‑11 kJ 8. Calculate ΔH for the reaction C2H4 (g) + H2 (g) → C2H6 (g), from the following data. C2H4 (g) + 3 O2 (g) → 2 CO2 (g) + 2 H2O (l) ΔH = -­‐‑1411. kJ/mole C2H6 (g) + 7/2 O2 (g) → 2 CO2 (g) + 3 H2O (l) ΔH = -­‐‑1560. kJ/mole H2 (g) + ½ O2 (g) → H2O (l) ΔH = -­‐‑285.8 kJ/mole -­‐‑136.8 kJ 9. Calculate ΔH for the reaction 4 NH3 (g) + 5 O2 (g) → 4 NO (g) + 6 H2O (g), from the following data. N2 (g) + O2 (g) → 2 NO (g) ΔH = -­‐‑180.5 kJ N2 (g) + 3 H2 (g) → 2 NH3 (g) ΔH = -­‐‑91.8 kJ 2 H2 (g) + O2 (g) → 2 H2O (g) ΔH = -­‐‑483.6 kJ -­‐‑1628.2 kJ I can… § state the sign of ΔH based on observation of warming or cooling of the surroundings. § correctly apply the terms exothermic and endothermic to situations where the surroundings are warming or cooling. Measuring Heat § state the units of heat capacity, specific heat, and molar heat capacity as well as the significance of each. § use calorimetry (q=mCΔT) to calculate heat changes during temperature changes. § calculate the heat transferred when two objects, at different temperatures, come into contact. Energy = Heat and Work § state the difference between work and heat energy. § state the difference between system and surroundings. § recognize the system and the surroundings in a chemical or physical system. § calculate the change in internal energy based on changes in heat absorbed by the system and work done by the system. § state that ΔH is a more general (and useful) measure of energy than ΔE and that ΔH = q when a reaction occurs at constant pressure. Chemical Work = Expanding Gases § relate physical work (w=F·d) and chemical work (w=-­‐‑P·ΔV). § calculate PV work done by an expanding gas. § state that no work is done in a constant volume situation such as a bomb calorimeter. Calculating ΔH -­‐‑-­‐‑ Hess’s Law § state the definition of a state function. § list examples of properties that are and are not state functions. § write the equation for the heat of formation of a substance. § state that the heat of formation of an element under standard conditions has a value of zero. § use Hess’s Law to calculate the energy of a chemical or physical change. C2H2(g) + 2 H2(g) -­‐‑> C2H6(g) Information about the substances involved in the reaction represented above is summarized in the following tables. Substance DH°f (kJ/mol) C2H2(g) 226.7 C2H6(g) -­‐‑84.7 (a) Write the equation for the heat of formation of C2H6(g) 2C (s) + 3H2 (g) → C2H6 (g) (b) Use the above information to determine the enthalpy of reaction for the equation given. -­‐‑311.4 kJ C6H5OH(s) + 7 O2(g) -­‐‑> 6 CO2(g) + 3 H2O(l) When a 2.000-­‐‑gram sample of pure phenol, C6H5OH(s), is completely burned according to the equation above, 64.98 kilojoules of heat is released. Use the information in the table below to answer the questions that follow. Standard Heat of Formation, Substance DH°f; at 25°C (kJ/mol) CO2(g) -­‐‑393.5 H2O(l) -­‐‑285.85 C6H5OH(s) ? (a) Calculate the molar heat of combustion of phenol in kilojoules per mole at 25°C. -­‐‑3058 kJ/mol (b) Calculate the standard heat of formation, DH°f, of phenol in kilojoules per mole at 25°C. -­‐‑160.6 kJ/mol 1. How many joules are equivalent to 37.7 cal? a) 9.01 J c) 1.51 J b) 4.184 J d) 158 J 2. The quantity of heat that is needed to raise the temperature of a sample of a substance 1.00 degree is called its a) heat capacity c) enthalpy b) specific heat d) kinetic energy 3. Equal masses of two substances, A & B, each absorb 25 Joules of energy. If the temperature of A increases by 4 degrees and the temperature of B increases by 8 degrees, one can say that a) the specific heat of A is double that of B. b) the specific heat of B is double that of A. c) the specific heat of B is negative. d) the specific heat of B is triple that of A. 4. If 25 J are required to change the temperature of 5.0 g of substance A by 2.0°C, what is the specific heat of substance A? a) 250 J/g°C c) 10. J/g°C b) 63 J/g°C d) 2.5 J/g°C 5. How much energy is required to change the temperature of 2.00 g aluminum from 20.0°C to 25.0°C? The specific heat of aluminum is 0.902 J/g°C. a) 2.3 J c) 0.36 J b) 9.0 J d) 0.090 J 6. Consider the thermal energy transfer during a chemical process. When heat is transferred to the system, the process is said to be _______ and the sign of ΔH is ________. a) exothermic, positive c) exothermic, negative b) endothermic, negative d) endothermic, positive
7. What is the ΔE for a system which has the following two steps: Step 1: The system absorbs 60 J of heat while 40 J of work are performed on it. Step 2: The system releases 30 J of heat while doing 70 J of work. a) 100 J c) 30 J b) 90 J d) zero 8. When two solutions react the container “feels hot.” Thus, a) the reaction is endothermic. b) the reaction is exothermic. c) the energy of the universe is increased. d) the energy of both the system and the surroundings is decreased. 9. The equation for the standard enthalpy of formation of N2O3 is a) N2O(g) + O2(g) → N2O3(g) b) N2O5(g) → N2O3(g) + O2(g) c) NO(g) + NO2(g) → N2O3(g) d) N2(g) + 3/2 O2(g) → N2O3(g) 10. For the general reaction 2 A + B2 → 2 AB, ΔH is +50.0 kJ. We can conclude that a) the reaction is endothermic. b) the surroundings absorb energy. c) the standard enthalpy of formation of AB is -­‐‑50.0 kJ. d) the molecule AB contains less energy than A or B2. 11. Calculate the enthalpy of combustion of C3H6 [C3H6(g) + 9/2O2(g) → 3CO2 + 3H2O] using the following: 3C(s) + 3H2(g) → C3H6(g) ΔH°= 53.3 kJ C(s) + O2(g) → CO2(g) ΔH°=-­‐‑394 kJ H2(g) + 1/2O2(g) → H2O(l) ΔH°=-­‐‑286 kJ a) -­‐‑1517 kJ c) -­‐‑626 kJ b) 1304 kJ d) -­‐‑2093 kJ 12. Which one of the following would have an enthalpy of formation value (ΔHf) of zero? a) H2O(g) c) H2O(l) b) O(g) d) O2(g) 13. Calculate the heat of vaporization of titanium (IV) chloride: TiCl4(l) → TiCl4(g) using the following enthalpies of reaction: Ti(s) + 2Cl2(g) → TiCl4(l) ΔH°=-­‐‑804.2 kJ TiCl4(g) → 2Cl2(g) + Ti(s) ΔH°= 763.2 kJ a) -­‐‑1567 kJ c) 1165 kJ b) -­‐‑783.7 kJ d) 41 kJ 14. 15. Calculate the enthalpy of reaction for D + F → G + M using the following equations and data: G + C → A + B ΔH° = +277 kJ C + F → A ΔH° = +303 kJ D → B + M ΔH° = -­‐‑158 kJ a) -­‐‑132 kJ c) +422 kJ b) -­‐‑422 kJ d) +132 kJ Calculate the standard enthalpy of the reaction for the process 3NO(g) → N2O(g) + NO2(g) using the standard enthalpies of formation (in kJ/mol): NO = 90; N2O = 82.1; NO2 = 34.0 a) -­‐‑153.9 kJ c) -­‐‑26.1 kJ b) 206 kJ d) 386 kJ 16. The standard molar enthalpy of combustion is -­‐‑1277.3 kJ for the combustion of ethanol. C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g) Calculate the standard molar enthalpy of formation for ethanol based on the following standard enthalpies of formation: ΔH°f CO2 = -­‐‑393.5 kJ/mol ΔH°f H2O = -­‐‑241.8 kJ/mol a) -­‐‑642.7 kJ/mol c) 235.1 kJ/mol b) -­‐‑235.1 kJ/mol d) 642.7 kJ/mol 17. Calculate the amount of heat needed to change 25.0 g ice at 0°C to water at 0°C. The heat of fusion of H2O = 333 J/g; a) 56.5 kJ c) 7.06 kJ b) 8.33 kJ d) 463 kJ 1. If the temperature of a 50.0-­‐‑gram block of aluminum increases by 10.9 K when heated by 500 Joules, calculate the a. heat capacity of the aluminum block 45.9 J/K b. specific heat of aluminum 0.917 J/g*K 2. Calculate the heat necessary to change the temperature of one kg of iron from 25°C to 1000°C. The specific heat of iron is 0.451 JK-­‐‑1g-­‐‑1. 440 kJ 3. If a 40 gram block of copper at 100°C is added to 100 grams of water at 25°C, calculate the final temperature assuming no heat is lost to the surroundings. The specific heat of copper is 0.385 JK-­‐‑
1g-­‐‑1 and the specific heat of water is 4.184 JK-­‐‑1g-­‐‑1. 27.7˚C 4. Calculate the amount of heat needed to melt 27.0 g of ice if the heat of fusion of ice is 6.009 kJ/mol. 9.003 kJ 5. If 27.0 grams of ice at 0°C is added to 123 grams of water at 100°C in an insulated container, calculate the final temperature. Assume that the specific heat of water is 4.184 JK-­‐‑1g-­‐‑1. 67.7˚C 6. A 50 gram block of an unknown metal alloy at 100°C is dropped into an insulated flask containing approximately 200 grams of ice. It was determined that 10.5 grams of the ice melted. What is the specific heat capacity of the unknown alloy? 0.70 J/g*K 7. If the enthalpy change for the combustion of propane is –2220 kJ/mole propane, what quantity of heat is released when 1 kg of propane is burned? C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l) ΔH-­‐‑2220 kJ -­‐‑50328 kJ 8. Using the following thermochemical data, calculate the molar heat of combustion, ΔH°combustion of methane, CH4: 2CH4(g) + 3O2(g) → 2CO(g) + 4H2O(l) ΔH° = -­‐‑1215 kJ 2C(s) + O2(g) → 2CO(g) ΔH° = -­‐‑221 kJ C(s) + O2(g) → CO2(g) ΔH° = -­‐‑394 kJ -­‐‑891 kJ/mol 9. Calculate the standard molar enthalpy of formation of methane from the data given in question 8, your answer to question 8, and the following: ΔH°f (H2O(l)) = -­‐‑286 kJ/mol -­‐‑75 kJ/mol 10. When ammonia is oxidized to nitrogen dioxide and water, the quantity of heat released equals 349 kJ per mol of ammonia: 2NH3(g) + 7/2O2(g) → 2NO2(g) + 3H2O(l) ΔH° = -­‐‑698 kJ Calculate the standard molar enthalpy of formation of ammonia if ΔH°f (H2O(l)) = -­‐‑286 kJ/mol ΔH°f (NO2(g)) = +33 kJ/mol -­‐‑47 kJ/mol 11. When 40 grams of ammonium nitrate is dissolved in 100 grams of water in a constant-­‐‑pressure coffee-­‐‑cup calorimeter, the temperature of the solution drops by 22.4°C. If the specific heat of the solution is 4.18 JK-­‐‑1g-­‐‑1, calculate the enthalpy of solution of ammonium nitrate. 26.2 kJ/mol