AP Chemistry 5: Thermochemistry & Thermodynamics

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1 Page 1 of 35 Name: AP Chemistry 5: Thermochemistry & Thermodynamics Class: Topics/ Daily Outline: Day A B Content TEXT CW HW 1 2/28 2/27 Measuring heat flow 6.1, /4 3/1 Hand warmer design /7 3/6 Finding H without experiment 6.3, /11 3/8 Spontaneity /13 3/12 Free energy /15 3/14 Free energy and chemical reactions /19 3/18 Unit Test Important Due Dates: o Hand Warmer Design, 3/13 (A Day) and 3/12 (B Day) For tutorials and additional resources: If you are absent, use this sheet to determine what you missed and collect the appropriate materials from your instructor. Get help from a friend, the link above, or the instructor. HW Assignments HW 1: Section 6.1 and 6.2. Cornell Notes Due 2/26 (A) & 2/27 (B) Read Sections 6.1 and 6.2 and complete Cornell Notes. OWL assignments will close at 11:45 PM on 3/19 HW 2: Heat and Enthalpy Due: 3/11 (A) & 3/8 (B) Complete through OWL. HW 3: Spontaneity Due: 3/13 (A) & 3/12 (B) Complete through OWL. HW 4: Review for Unit Test Due: 3/19 (A) & 3/18 (B) Complete the review using your test prep book (Topic 19), hand in during class.

2 Drills Page 2 of 35 Date Outcome Drill Date Outcome Drill

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5 CW 1: Energy and Heat Flow Page 5 of 35 Important Terms and Concepts Energy The capacity to do work or to produce heat. o Heat, q: Energy transfer between two objects due to a difference in temperature. q = mc T o Work, w: Force acting over a distance. In the diagram below, the energy due to the position of A is used to push the ball up the smaller hill (work). For a gas, work can be found from the product of pressure and volume change. w = P V Law of Conservation of Energy Energy can be converted from one form to another but cannot be created nor destroyed. Potential Energy, PE Energy due to position, arrangement, or composition o Water behind a dam: PE is converted to work as water runs turbines to make electricity. o Attractive and repulsive forces: The energy change of a chemical reaction results from differences in attractive forces between the nuclei and electrons in the reactants and products. Kinetic Energy, KE Energy due to motion, depends on mass and velocity. KE = 1 2 mv 2 State Function/ Property A property of the system that depends only on its present state it doesn t matter how you arrived at the present state. Altitude is a state function; distance is not PE is a state function

6 Page 6 of 35 The Universe We divide the universe into two parts: o System: What we are studying, often the reactants and products of a chemical reaction. o Surroundings: Everything else. Endothermic A system, reaction, or process that absorbs heat from the surroundings. Exothermic A system, reaction, or process that releases heat to the surroundings. First Law of Thermodynamics The total energy of the universe is constant.

7 Page 7 of 35 Internal Energy The sum of all potential and kinetic energy of all the particles in the system. May be changed through heat, work, or both. E = q + w Enthalpy The sum of a system s internal energy (E) and the product of its pressure (P) and volume (V). At constant pressure, the enthalpy change ( H) is equal to the heat change (q). Specific heat, c The amount of energy required to raise one gram of a substance by one degree (J/g C) or the amount of energy required to raise one mole of a substance by one degree (J/mol C). Calorimeter Insulated device designed to measure heat flow during a chemical reaction or process. o Calorimeter constant: The amount of heat absorbed per degree by the device itself; used as a correction factor, units of J/ C. o Constant pressure calorimetry: Measurement of heat flow under constant pressure (usually of the atmosphere). o Constant volume calorimetry: Measurement of heat flow under constant volume (achieved using a bomb calorimeter). Because volume is constant, no PV work is done. A constant pressure calorimeter; open to the atmosphere A constant volume calorimeter (or bomb calorimeter). Typically used to study combustion reactions: a sample is combusted via electrical ignition wires in the inner bomb; this energy is transferred to the water.

8 Problems and Questions Complete the following on a separate sheet of paper. Page 8 of Calculate the change in internal energy ( E) for a system undergoing an endothermic process in which 15.6 kj of heat flows and 1.4 kj of work is done on the system. 2. Calculate the work associated with the expansion of a gas from 46 L to 64 L at a constant external pressure of 15 atm. 3. The volume of a hot air balloon changes from L to L by the addition of J of energy as heat. Assuming the pressure is constant at 1.0 atm, calculate the change in internal energy ( E) for this process. 4. If the internal energy of a thermodynamic system is increased by 300. J while 75 J of expansion work is done, how much heat was transferred and in which direction? 5. When 1 mol of methane (CH 4 ) is burned at constant pressure, 890 kj of energy is released as heat. Calculate the change in enthalpy ( H) if 5.8 g of methane is burned at constant temperature. 6. When Ba(NO 3 ) 2 solution is reacted with Na 2 SO 4 solution, the white solid BaSO 4 forms. Use the data below to determine the enthalpy change per mole of BaSO 4 formed. Ba(NO 3 ) 2 solution 1.00 L of 1.00 M Ba(NO 3 ) 2 Na 2 SO 4 solution 1.00 L of 1.00 M Na 2 SO 4 T initial (both solutions) 25.0 C T final (final solution, mixed) Specific heat of final solution Density of final solution 28.1 C J/g C (dilute solution = same as H 2O) 1.0 g/ml (dilute solution = same as H2O) 7. Are the following processes endothermic or exothermic? a. When solid KBr is dissolved in water, the solution gets colder. b. Propane is burned on a grill c. Water condensing on a cold pipe d. CO 2 (g) CO 2 (s) e. F 2 (g) 2F (g) 8. The combustion of g of benzoic acid increases the temperature of water in a bomb calorimeter by 2.54 C. If the energy released by combustion of benzoic acid is kj/g, calculate the heat capacity of the calorimeter in kj/ C. 9. A g sample of vanillin (C 8 H 8 O 3 ) is then burned in the same calorimeter, causing the temperature of the water inside to increase by 3.25 C. What is the energy of combustion per gram of vanillin? Per mole of vanillin?

9 CW 2: Designing a Hand Warmer Page 9 of 35 Challenge You will use chemistry to design an effective, safe, environmentally benign, and inexpensive hand warmer. The ideal hand warmer increases in temperature by 20 C (but no more) as quickly as possible, uses about 50 grams of water, costs as little as possible to make, and uses chemicals that are as safe and environmentally friendly as possible. You will carry out an experiment to determine which substances, in what amounts, to use to make a hand warmer that meets these criteria. Background People put hand warmers inside their gloves on cold days to keep their fingers warm. They are popular with people who work outside in winter or engage in winter sports. One type of hand warmer consists of a packet that contains water in one section and a soluble substance in another section. When the packet is squeezed the water and the soluble substance are mixed, the solid dissolves and the packet becomes warm. Breaking chemical bonds or particulate attractions absorbs energy from the surroundings, while forming new bonds and particulate attractions releases energy to the surroundings. When an ionic solid dissolves in water, ionic bonds between cations and anions and hydrogen bonds between water molecules are broken. New attractions between water molecules and the cations and anions are formed. The amount of energy involved to break these bonds and form new ones depends on the chemical properties of the anions and cations. Therefore, when some ionic solids dissolve, more energy is required to break the bonds/ attractions than is released in forming the new water-ion attractions, and the overall process absorbs energy in the form of heat. If this occurs at constant pressure, the change in enthalpy is the same as the heat change (q), and the process is endothermic. When other ionic compounds dissolve, bond making releases more energy than bond breaking absorbs, and the process releases heat and is exothermic. In this experiment, you will collect data to calculate the enthalpy of solution ( H soln, kj/mol solute) occurring in aqueous solution. The data necessary to calculate the heat of solution can be obtained using a calorimeter, a device designed to isolate the system and surroundings from the rest of the universe so that heat flow may be studied. This is accomplished by insulating the calorimeter to prevent heat loss.

10 In the calorimeter shown, a chemical reaction (system) occurs in water (surroundings). According to the law of conservation of energy, energy cannot be created or destroyed, only changed between forms or transferred from one system to another. The thermometer measures the temperature of the water, which can be used to determine the heat absorbed or released by the reaction. H water = H reaction Page 10 of 35 Recall that at constant pressure, the enthalpy change ( H) is equal to the heat change (q). q water = H reaction m water (4.184 J g ) T water = H reaction During the experiment, the calorimeter itself may absorb (or release) a small amount of heat. The amount of heat gain (or loss) by the calorimeter can be used as a correction factor in heat calculations, known as the calorimeter constant. The calorimeter constant is determined by mixing hot and cold water in the calorimeter. When hot and cold water are mixed, the hot water transfers some of its heat to the cool water. The law of conservation of energy dictates that the amount of heat lost by the hot water, q hot, is equal to the heat gain of the cool water, q cold, but opposite in sign. Using the mass of water, the temperature change, and the specific heat of water (4.184 J/g C), q hot and q cold may be determined. q hot = q cold m hot (4.184 J g ) T hot = m cold (4.184 J g ) T cold The difference in these two numbers is due to heat exchange with the calorimeter. The calorimeter constant (C) is found by dividing the heat change of the calorimeter by the temperature change of the calorimeter. C = enthalpy change of the calorimeter temperature change if the calorimeter

11 Lab Notebook and Report Page 11 of 35 In the lab notebook: PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB 1. Introduction: Explain theory and applications behind the lab, as well as any chemical equations, graphs, diagrams, or mathematical equations that will be required. o Explain the criteria for the hand warmer design o Describe the hand warmer packet, why mixing the sections releases heat o Explain calorimetry, why and how to find the calorimeter constant 2. Safety Data Sheets: Look up the SDS for all chemicals used during the lab use the table below to list major hazards (typically in the Hazards Identification section). Chemical Eye Irritant Skin Irritant Respiratory Irritant Other Info NH 4 NO 3 Na 2 CO 3 CaCl 2 NaCl LiCl NaC 2 H 3 O 2 Summary of Safety Precautions: 3. Materials and Methods: Complete the corresponding classwork in the unit packet. Write a procedure listing all chemicals and equipment. Do not worry about sizes of glassware, concentrations/ amounts of chemicals, or detailed LoggerPro instructions. o Write a procedure for finding the calorimeter constant o Write a separate procedure for determining H soln PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB PRE-LAB 4. Procedural Notes: Record details of the procedure, exact sizes of glassware, concentrations or amounts of chemicals, instructions for using LoggerPro, and any other notes provided by your instructor or group members. 5. Data Collection: Neatly record any measurements, written with the correct precision based on the equipment, detailed observations and possible sources of error. o Calorimeter Constant Data: mass of hot water, mass of cold water, initial temperature of hot and cold water, final temperature of the mixture o Hand Warmer Data: solid, cost per gram of solid, molar mass of solid, exact mass of water, initial water temperature, final water temperature 6. Calculations: Show one handwritten worked example for each major required calculation. Include the unrounded answer and the answers with correct sig figs. You do not need to show sample calculations for finding averages or standard deviations. o Calculate your calorimeter constant This should have units of J/ C. o Calculate the heat of solution for each of your solids, taking the calorimeter constant into consideration. This should in in units of kj/mol of solid. Report this value to the class data table.

12 Page 12 of 35 Typed and printed: 7. Excel Analysis: Access the class set of data using Complete all required calculations. Copy and paste your spread sheet and any required graphs into your typed document. o Find the average heat of solution and standard deviation for each solid. o Calculate the percent error in the heat of solution actual values will be supplied in the class set of data. 8. Conclusion: Include the following in narrative form: o Discussion of what you did, including the purpose/ goal of the experiment. Describe any changes to the procedure from what you wrote in your lab notebook. o Based on the heat of solution data and hand warmer criteria, describe your hand warmer design. Be specific with amounts and include calculations for the mass of solid required. o Comment on error, both inherit error (due to precision of equipment) and experimental/ human error. Remember that the lab would have been better if we did it correctly or human errors were made are not discussion of error. Errors should be clearly described along with how they impacted the calculations in the data analysis. o Describe refinements or future experiments. What can we study next and why? 9. Cover Sheet: Create a cover sheet using the exact table below. Place cover sheet on top of items 4 to 8 from above. Report Title: Hand Warmer Design My Name: Partners: Class Period: Due Date: 3/11 (A) & 3/12 (B)

13 CW 3: Finding H without Experiment Page 13 of 35 For a reaction studied under constant pressure, we can use calorimetry to find the enthalpy change. Sometimes it is difficult to use this process. The reaction may be too slow, or it may be an intermediate step in a series of reactions. These pen and paper methods provide another way to find ΔH. Hess Law Hess s law of heat summation allows you to determine the heat of reaction indirectly using known heats of reaction of two or more thermochemical equations. This is possible because enthalpy is a state function. Consider the conversion of diamond to graphite. This reaction is too slow to measure directly. C(s, diamond) C(s, graphite) ΔH =? (a) The enthalpy change is known for the following reactions: C(s, diamond) + O 2 (g) CO 2 (g) ΔH = kj (b) C(s, graphite) + O 2 (g) CO 2 (g) ΔH = kj (c) Write reaction (c) in reverse, starting with products and creating reactants. Because you reversed the direction of the reaction, you must also reverse the sign of the ΔH value. ΔH = Write reaction (b) and the reversed reaction (d) in the boxes below. Combine and cancel reactions to get reaction (a) and find the value of ΔH. (b) ΔH = (d) ΔH = (a) ΔH = (d) Normal algebraic relationships may be applied to chemical equations. For example, 1 mole of A left of the arrow cancels 1 mole of A right of the arrow, leaving behind 2 moles of A. 3A + B C + D C + D 1A + E 2A + B E You may multiply a reaction by any whole positive number, so long as you multiply the ΔH by the same whole number. A + B C ΔH = 10 kj 2A + 2B 2C ΔH = 20 kj

14 1. Given the following data Page 14 of 35 Calculate H for the reaction 2. Given the following data: 2ClF(g) + O 2 (g) Cl 2 O(g) + F 2 O(g) 2ClF 3 (g) + 2O 2 (g) Cl 2 O(g) + 3F 2 O(g) 2F 2 (g) + O 2 (g) 2F 2 O(g) Calculate the H for the reaction: ClF(g) + F 2 (g) ClF 3 (g) H = kj H = kj H = 43.4 kj

15 Page 15 of 35 Standard Enthalpy of Formation The standard enthalpy of formation ( H f ) of a compound is defined as the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states (denoted by the degree sign). The standard state for a substance is a precisely defined reference state. Because enthalpy is a state function, the difference in these standard reference states can be used to determine the enthalpy of a reaction. The change in enthalpy for a reaction can be calculated by subtracting the enthalpies of formation of the reactants from the enthalpies of formation of products. Remember to multiply the enthalpies of formation by their molar coefficients from the balanced chemical equation. An important note is that H f for an element in its standard state (whatever state it would be at 1 atm and 25 C) is zero. H reaction = n p H f (products) n r H f (reactants) Here are some standard enthalpies for formation for several compounds. A more comprehensive table is found in Appendix 4 of your text. Compound H f (kj/mol) NH 3 (g) 46 NO 2 (g) 34 H 2 O(l) 286 Al 2 O 3 (s) 1676 Fe 2 O 3 (s) 826 CO 2 (g) 394 CH 3 OH(l) 239 C 8 H 18 (l) Calculate the standard enthalpy change for the reaction that occurs when ammonia is burned in air to form nitrogen dioxide and water. This is the first step in the manufacture of nitric acid. 4NH 3 (g) + 7O 2 (g) 4NO 2 (g) + 6H 2 O(l)

16 4. Calculate the standard enthalpy change for the thermite reaction. 2Al(s) + Fe 2 O 3 (s) Al 2 O 3 (s) + 2Fe(s) Page 16 of Until recently, methanol (CH 3 OH) was used as fuel in higher performance engines in race cars. Compare the standard enthalpy of combustion per gram of methanol and per gram of gasoline. Gasoline is really a mixture of compounds but assume for this problem that gasoline is pure liquid octane (C 8 H 18 ).

17 CW 4: Spontaneity Page 17 of 35 Many physical and chemical processes proceed naturally in one direction, but not the other. Consider the reaction between zinc and hydrochloric acid: Zn(s) + HCl(aq) ZnCl 2 (aq) + H 2 (g) When zinc is added to the acid, it begins reacting immediately, forming H 2 gas bubbles and ions in solution. We do not observe the reverse reaction it would seem very odd for the bubbles to form at the surface of the solution and sink down as the metal zinc strip reformed. Reactions are spontaneous in the direction in which they naturally occur. What factor(s) determine the direction in which reactions are spontaneous? A Ball Rolls Downhill 1. Considering the diagram above: a. Which is the lower energy state: the ball at the top of the hill or at the bottom? b. Is the change in potential energy for this process positive, negative or zero? c. Why doesn t the ball roll uphill?

18 Solid NaCl is Formed from Gaseous Ions Page 18 of Based on the diagram: a. Which is the lower energy state: 1 mol of solid NaCl or 1 mol of Na + and Cl ions? b. Is the H for this process Na + (g) + Cl (g) NaCl(s) positive, negative, or zero? c. Why doesn t a salt crystal suddenly become gaseous sodium and chloride ions? 3. Write a chemical equation that describes the freezing of water, and indicate whether ΔH for the process is positive, negative, or zero. a. Under what temperature conditions does this process naturally occur? b. Is it possible to determine whether a process will occur naturally solely by examining the sign of ΔH for the process? Explain. c. What other factor(s) besides the sign of ΔH must be considered to determine if a process will occur spontaneous under a given set of conditions?

19 Page 19 of 35 Disorder Drives the Universe Every developers of thermodynamics thought that a process would be spontaneous if it was exothermic. This explanation is not complete, as it cannot account for the melting of ice at room temperature, an endothermic process. If this is the answer, why does the melting of ice occur spontaneously at room temperature, an endothermic process? The characteristic common to all spontaneous processes is an increase in entropy (S). Entropy can be viewed as a measure of molecular randomness or disorder. The natural order is from low entropy to higher entropy (more disorder). Entropy is a thermodynamic function that describes the number of arrangements that are available to a system and is closely related to probability. We can use this to explain why the gas below spontaneously expands to fill both bulbs. 4. Considering the arrangements above: a. Which arrangement would you expect to find in the real world? b. Which arrangement has higher entropy? c. How do Questions 4a and 4b relate to each other and to probability? Explain.

20 Arrangement I (4 in Left) Microstates # Available States Page 20 of 35 Probability 1 1/16 = II (3 in Left 1 in Right) 4 III (2 in Left 2 in Right) 6 IV (1 in Left 3 in Right) 4 V (4 in Right) 1 5. The diagram above shows all possible arrangements (states) of four gas particles in a two-bulb flask. a. Which arrangement has the highest entropy? I II III IV V b. Calculate the probability of encountering each possible arrangement. Which of these states is most probable? c. If two more particles are added, how would the probability for encountering arrangement I (all on the left) change? Arrangement III (evenly distributed)? d. Return to Question 4. Why is B more likely? Explain.

21 Page 21 of 35 We have seen that processes are spontaneous when they result in an increase in disorder. Nature always moves towards the most probable state available to it. This leads to the second law of thermodynamics: the driving force for a spontaneous process is an increase in the entropy of the universe. In other words, the entropy of the universe is increasing. We can divide the universe into system and surroundings: S univ = S sys + S surr o If ΔS unvi is positive, the entropy of the universe increases, and the process is spontaneous in the direction written. o If ΔS unvi is negative, the process will be spontaneous in the opposite direction. o If ΔS unvi is zero, the process has no tendency to occur, and the system is at equilibrium. 6. In a living cell, large molecules are assembled from simple ones. Is this process consistent with the second law of thermodynamics? Molecules in a Box 7. Which system has the greatest entropy? 8. Suppose that at some temperature the process above starts with A and ends with C. a. Is S A > S C or is S C > S A? b. Is ΔS for this naturally occurring process positive or negative? c. Suppose that at some different temperature, the process starts with C and ends with A. Is ΔS for this naturally occurring process positive or negative?

22 Page 22 of 35 The following generalizations are useful to find the change in entropy (ΔS) for a process: o As the number of particles increases, disorder increases. (+ΔS) o As the volume in which particles can move increases, disorder increases. (+ΔS) o As the temperature (KE) of particles increases, disorder increases. (+ΔS) 9. For each of the following, is ΔS positive or negative? Explain. a. C 4 H 8 (g) 2C 2 H 4 (g) b. C 4 H 8 (g) C 4 H 8 (s) c. C 4 H 8 (g, 298K, 1 atm) C 4 H 8 (g, 298K, 0.5 atm) d. C 4 H 8 (g, 298K, 1 atm) C 4 H 8 (g, 398K, 1 atm) Sodium Chloride Dissolves in Water 10.Considering the dissolving of solid NaCl in water: a. Is this process endothermic or exothermic? b. Does entropy increase or decrease? 11.How does Question 9 relate to spontaneity for this process?

23 The Relationship Between ΔH, ΔS, and T ΔH ΔS Higher T? Lower T? + + Yes No + No No + Yes Yes No Yes Page 23 of In the table above: a. Which row above corresponds to the melting of ice? b. Which row above corresponds to the freezing of ice? c. Is the data consistent with the statement that exothermic reactions tend to be spontaneous? Explain. d. For what values of ΔS do chemical reactions occur spontaneously? How does this relate to the second law of thermodynamics? 13.Under what conditions will an endothermic reaction be spontaneous? 14.For the following, indicate the driving forces that make the process spontaneous. Process Entropy Enthalpy Both When NH 4 NO 3 (s) dissolves in water the temperature of the solution decreases. When concentrated sulfuric acid is added to water, the acid solution is diluted and the temperature of the solution increases.

24 Page 24 of 35 Recall that the change entropy of the universe is found from the change in entropy of the system and surroundings: S univ = S sys + S surr Case ΔS sys ΔS surr ΔS unvi A B C +? D +? 15.Considering each case: a. Why is case A spontaneous? Why is case B nonspontaneous? Explain. b. Can you determine if cases C and D are spontaneous or not? Explain. 16.The freezing of water can be expressed as: H 2 O(l) H 2 O(g) a. Is the freezing of water endothermic or exothermic? b. Is the change in entropy for this system (ΔS sys ) positive or negative? c. What values of ΔS surr allow this process to occur spontaneously?

25 Page 25 of 35 Entropy changes in the surroundings (ΔS surr ) primary depend on heat flow: o The sign of ΔS surr depends on the direction of heat flow. If heat flows from the system into the surroundings (exothermic), the heat increases the random motions of particles, increasing ΔS surr. If heat flows out of the surroundings into the system (endothermic), the ΔS surr decreases. o The magnitude of ΔS surr depends on the amount of heat transferred and temperature: Exothermic process: S surr = + quantity of heat (J) temperature (K) Endothermic process: S surr = quantity of heat (J) temperature (K) At constant pressure, the quantity of heat transferred is equal to the change in enthalpy: S surr = H T The minus sign is necessary because the sign of H is determined with respect to the system, and this equation expresses a property of the surroundings. 17.If a system is exothermic, does the temperature matter in determining the sign of ΔS surr? What if the system is endothermic? 18.In the metallurgy of antimony, the pure metal is recovered via different redox reactions, depending on the composition of the ore. In sulfide ores, iron is used to reduce antimony: Sb 2 S 3 (s) + 3Fe(s) 2Sb(s) + 3FeS(s) In oxide ores, carbon is used to reduce antimony: Sb 4 O 6 (s) + 6C(s) 4Sb(s) + 6CO(g) ΔH = 125 kj ΔH = 778 kj Calculate the ΔS surr for each reaction at 25 C and 1 atm.

26 CW 5: Free Energy Page 26 of 35 Another Way to Predict Spontaneity So far, we have used the change in entropy to predict if a process occurs naturally (spontaneously). Another thermodynamic function, free energy (G), is useful in dealing with the temperature dependence of spontaneity. For a process that occurs at constant temperature, the change in free energy is found by: Where ΔG: Change in free energy G = H T S ΔH: Change in enthalpy T: Kelvin temperature ΔS: Change in entropy A process (at constant T and P) is spontaneous for all values of T that result in a negative free energy. These conditions of constant T and P are encountered often by chemists, so this equation is very useful. Melting Ice The data below is for the melting of ice at three different temperatures: T ( C) T (K) ΔH (J/mol) ΔS (J/mol K) TΔS (J/mol) ΔG = ΔH TΔS Note: the degree sign ( ) indicates that all substances are in their standard states (CW 3). 1. Complete the chart to find ΔG. a. Based on your calculations, at which temperature is this process spontaneous? b. Does this match observations you have about melting ice? Explain.

27 Page 27 of Complete the chart below by indicating if each case below is spontaneous at high and low temperatures. Case ΔS ΔH Low Temp? High Temp? A + B + + C D + 3. Use the free energy equation to explain how a process can be endothermic and spontaneous (case B) and exothermic and nonspontaneous (case C). 4. Which of the following processes are spontaneous under the given conditions? a. ΔH = +25 kj, ΔS = +5.0 J/K, T = 300. K b. ΔH = +25 kj, ΔS = J/K, T = 300. K c. ΔH = 10. kj, ΔS = +5.0 J/K, T = 298 K d. ΔH = 10. kj, ΔS = 40. J/K, T = 200. K 5. At what temperatures will the following processes be spontaneous? a. ΔH = 18 kj and ΔS = 60 J/K b. ΔH = +18 kj and ΔS = +60 J/K c. ΔH = +18 kj and ΔS = 60 J/K d. ΔH = 18 kj and ΔS = +60 J/K

28 You ve probably been hearing like dissolves like since the 6 th grade now you re a junior or senior in high school! It took that long! Entropy Changes in Aqueous Solutions Recall the concept of like dissolves like. This implies that an ionic solute (such as NaCl) should dissolve in a polar solvent (such as H 2 O). This process can be described as a series of steps: o Breaking solute-solute attractions (endothermic), such as bond energy o Breaking solvent-solvent attractions (endothermic), such as H bonding o Forming solvent-solute attractions (exothermic), in solvation Page 28 of Thermodynamic data for several solution processes are given below. Solute Solvent ΔH 1 ΔH 2 ΔH 3 ΔH soln ΔS T ΔG Polar Polar Large, + Large, + Large, + or Nonpolar Nonpolar Small, + Small, + Small, + or Nonpolar Polar Small, + Large, + Small, + or Polar Nonpolar Large, + Small, + Small, + or a. Can you easily predict ΔH soln? Why? b. Which processes are spontaneous? c. Why does NaCl dissolve in water despite having ΔH soln = 3 kj/mol? d. In terms of ΔS, use the diagram below to explain why a nonpolar solute won t dissolve in a polar solvent.

29 7. Thermodynamic data is given below for the dissolving process of several salts. Process ΔS ΔH Soluble in H 2 O? ΔG KCl K + (aq) + Cl (aq) + + LiF Li + (aq) + F (aq) + CaS Ca 2+ (aq) + S 2 (aq) Page 29 of 35 a. Use solubility rules to identify each salt as soluble or insoluble in water. b. Based on your previous answer, will ΔG will be + or for each process? c. Explain why some salts are soluble despite negative entropy d. How does increasing the temperature effect the value of ΔG? Does this increase or decrease the solubility of a salt? Entropy Changes in Chemical Reactions Because entropy is a state function of a system, the change in entropy for a given chemical reaction can be calculated by taking the difference between the standard entropy values of products and reactants: S reaction = n p S products n r S reactants 8. Calculate ΔS at 25 C for the reaction using the given standard entropy values. 2NiS(s) + 3O 2 (g) 2SO 2 (g) + 2NiO(s) Substance S (J/molK) SO 2 (g) 248 NiO(s) 38 O 2 (g) 205 NiS(s) 53

30 Finding ΔG for Chemical Reactions Page 30 of 35 Method 1: Free Energy Equation 9. Consider the reaction 2SO 2 (g) + O 2 (g) 2SO 3 (g) carried out at 25 C and 1 atm. Calculate ΔH, ΔS, ΔG using the following data: Substance H f (kj/mol) S (J/mol K) SO 2 (g) SO 3 (g) O 2 (g) Method 2: Like Hess Law Free energy is a state function (like enthalpy), allowing us to find ΔG much like Hess law. 10. Determine the free energy change for the reaction: 2CO(g) + O 2 (g) 2CO 2 (g) from the following data: 2CH 4 (g) + 3O 2 (g) 2CO(g) + 4H 2 O(g) ΔG = 1088 kj CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(g) ΔG = 801 kj Method 3: Standard Free Energy of Formation G reaction = n p G f (products) n r G f (reactants) 11.Methanol is a high-octane fuel used in high-performance racing engines. Calculate ΔG for the reaction: 2CH 3 OH(g) + 3O 2 (g) 2CO 2 (g) + 4H 2 O(g) Given the following free energies of formation: Substance ΔG f (kj/mol) CH 3 OH(g) 163 O 2 (g) 0 CO 2 (g) 394 H 2 O(g) 229

31 CW 6: Thermodynamics and Equilibrium Page 31 of 35 Free Energy and Equilibrium We have seen that when the products are at a lower enthalpy than the reactants ( ΔH, exothermic), a chemical reaction is energetically favored. We have also seen that when the products are more disordered than the reactants (+ΔS ), a chemical reaction is entropically favored. Reaction ΔH (kj/mol) Enthalpy favorable? ΔS (J/mol K) NaCl(s) Na + (aq) + Cl (aq) NH 4 NO 3 (s) NH 4+ (aq) + NO 3 (aq) Zn(s) + Cu 2+ (aq) Cu(s) + Zn 2+ (aq) Cl (aq) + Br 2 (l) 2Br (aq) + Cl 2 (g) CH 3 COOH CH 3 COO (aq) + H + (aq) Entropy favorable? 1. For each of the reactions in the table a. indicate if the reaction is favorable with respect to the change in enthalpy and the change in entropy. Enter Y or N into the table. b. Can you predict, from the given data, which processes are spontaneous? c. Are there any reactions where both the enthalpy and entropy factors are favorable? d. Are there any reactions where both the enthalpy and entropy factors are favorable?

32 Page 32 of 35 Reaction ΔH (kj/mol) T ΔS ΔH T ΔS (kj/mol) (kj/mol) K NaCl(s) Na + (aq) + Cl (aq) NH 4 NO 3 (s) NH 4+ (aq) + NO 3 (aq) Zn(s) + Cu 2+ (aq) Cu(s) + Zn 2+ (aq) Cl (aq) + Br 2 (l) 2Br (aq) + Cl 2 (g) CH 3 COOH CH 3 COO (aq) + H + (aq) Considering the thermodynamic and equilibrium data above a. Which reactions appear to be favorable (K>1)? b. Which reactions appear to be unfavorable (K<1)? c. Which factor is used to predict if the reaction is favorable or unfavorable? i. ΔH ii. T ΔS iii. ΔH T ΔS d. When K>1, is the factor you identified positive or negative? e. When K<1, is the factor you identified positive or negative? f. What is the quantitative relationship between the factor you identified and the magnitude of K? 4. Without referring to tables to calculate ΔH and ΔS, predict whether the equilibrium constant at room temperature for the following reaction will be greater than, less than, or equal to 1. Explain your reasoning. 5. For what combinations of ΔH and ΔS will a chemical reaction always have K>1? Always have K<1?

33 Page 33 of 35 The Meaning of ΔG for a Chemical Reaction We have learned that a system at constant temperature and pressure will proceed spontaneously in the direction that lowers its free energy. This is why reactions proceed until they reach equilibrium. The equilibrium position represents the lowest free energy value available to a reaction system. This leads to the following equations: For a reaction NOT at equilibrium: G = G + RTln(Q) For a reaction at equilibrium: G = RTln(K) (This can also be used to determine the change in free energy of a reaction at nonstandard temperatures and pressures) (When a reaction is at equilibrium, ΔG = 0, and Q = K. Substitute this into the previous equation and solve to get this one.) Where: o ΔG = Change in free energy at specified (nonstandard state) conditions o ΔG = Change in free energy at standard state conditions o R = Gas law constant = J/mol K o T = Kelvin temperature o Q = Reaction quotient o K = Equilibrium constant Recall that we can write these for both pressure and concentration. 1. One method for synthesizing methanol involves reacting carbon monoxide and hydrogen gases: CO(g) + 2H 2 (g) CH 3 OH(l) a. Calculate ΔG at 25 C for this reaction where carbon monoxide at 5.0 atm is mixed with hydrogen gas at 3.0 atm to produce liquid methanol. (HINT: Find ΔG f using Appendix 4) b. Based on the value for G, would you expect this reaction to be more spontaneous at high or low pressures? c. How does this relate to Le Châtelier s principle for the reaction above?

34 Page 34 of Consider the ammonia synthesis reaction: N 2 (g) + 3H 2 (g) 2NH 3 (g) Where ΔG = 33.3 kj/mole of N 2 consumed at 25 C. For each of the following mixtures of reactants and products at 25 C, predict the direction in which the system will shift to reach equilibrium. a. P NH3 = 1.00 atm, P N2 = 1.47 atm, P H2 = atm b. P NH3 = 1.00 atm, P N2 = 1.00 atm, P H2 = 1.00 atm 3. The overall reaction for the corrosion (rusting) of iron by oxygen is: 4Fe(s) + 3O 2 (g) 2Fe 2 O 3 (s) Using the following data, calculate the equilibrium constant for this reaction at 25 C. Substance ΔH f (kj/mol) S (J/mol K) Fe 2 O 3 (s) Fe(s) 0 27 O 2 (g) 0 205

35 Reference Materials Page 35 of 35 Polyatomic Ions Ion Name Ion Name H 3 O + hydronium 2 CrO 4 chromate 2+ Hg 2 dimercury (I) 2 Cr 2 O 7 dichromate + NH 4 ammonium MnO 4 permanganate C 2 H 3 O 2 NO acetate 2 nitrite CH 3 COO NO 3 nitrate 2 C 2 O 4 oxalate 2 O 2 peroxide 2 CO 3 carbonate OH hydroxide HCO 3 hydrogen (bi)carbonate CN cyanide 3 PO 4 Phosphate SCN thiocyanate ClO hypochlorite 2 SO 3 sulfite ClO 2 chlorite 2 SO 4 sulfate ClO 3 chlorate HSO 4 hydrogen sulfate ClO 4 perchlorate 2 S 2 O 3 thiosulfate Solubility Guidelines Ions that form Soluble Exceptions Compounds Group 1 ions (Li +, -- Na +, etc.) Ammonium (NH ) Nitrate (NO 3 ) -- Acetate (C 2 H 3 O 2 or CH 3 COOH) Hydrogen carbonate (HCO 3 ) Chlorate (ClO 3 ) -- Perchlorate (ClO 4 ) Halides (Cl, Br, I ) Sulfates (SO 4 2 ) -- When combined with Ag +, Pb 2+ 2+, Hg 2 When combined with Ag +, Ca 2+, Sr 2+, Ba 2+, Pb 2+ Ions that form Insoluble Compounds Carbonate (CO 2 3 ) Chromate (CrO 2 4 ) Phosphate (PO 3 4 ) Sulfide (S 2 ) Hydroxide (OH ) Exceptions When combined with Group + 1 ions or NH 4 When combined with Group 1 ions, Ca 2+, Mg 2+ + or NH 4 When combined with Group + 1 ions or NH 4 When combined with Group + 1 ions or NH 4 When combined with Group 1 ions, Ca 2+, Ba 2+, Sr 2+ + or NH 4