Chemistry work sheet

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Adapted from PLTL “The energies of our system will decay, the glory of the sun will be dimmed, and the earth, tideless and inert, will no longer tolerate the race which has for a moment disturbed its solitude. Man will go down into the pit, and all his thoughts will perish.” A.J. Balfour Enthalpy, Entropy, and Free Energy Calculations Can we predict if a reaction will occur? This is a question that can be answered by applying the principles of chemical thermodynamics, the study of the energy relationships associated with chemical reactions. The term we use to label a reaction that proceeds without continual energy input is spontaneous. A spontaneous reaction is one that will occur all by itself, once it has been given a small amount of energy so that it can get started. The burning of paper, for example, is a spontaneous reaction. Once you add a little bit of energy, like the heat from a match, the paper continues to burn without any outside help, until there is no more paper to burn. In contrast, a nonspontaneous reaction is one that will not proceed unless an outside source of energy is used. An example of a nonspontaneous reaction is the decomposition of water into hydrogen and oxygen. If we add energy to water, the water may begin to decompose (if the amount of energy is great enough) but the decomposition will stop as soon as the energy source is cut off. In this unit we will focus on performing calculations that will allow us to predict the spontaneity of chemical reactions. Free Energy J. Willard Gibbs (1839–1903) can be considered as one of the founding fathers of the field of chemical thermodynamics. He introduced a quantity known as Gibbs free energy, G, that represents the amount of energy available in a chemical system that can do useful work. If we are to consider a chemical change, we are interested in the change of free energy, or ∆G. Thus ∆G is a measure of the amount of energy in a chemical change that is free to do work on another physical or chemical system. The most notable aspect of the ∆G concept is that its sign allows us to predict the spontaneity of a chemical reaction that occurs at a constant temperature and pressure: 1. If ∆G is negative, the reaction is spontaneous. 2. If ∆G is positive, the reaction is nonspontaneous. 3. If ∆G is zero, the reacting system is at equilibrium, and there will be no change in the reaction on the macroscopic level. Enthalpy, Entropy, and Free Energy Calculations Standard state values of ∆G, symbolized as ∆G°, are commonly found in tables of thermodynamic quantities. Recall that the thermodynamic standard state conditions are 25°C, 1 atm pressure for gases, and 1 M concentrations for solutions. Calculation of ∆G° for a reaction is given by ∆G° = Σn∆Gf°products – Σn∆Gf°reactants (Eq.1) where ∆Gf° is the standard free energy of formation; that is, the free energy change that occurs when one mole of a compound is formed from elements in their standard states. Note the similarity of this equation to the equation used to calculate ∆H° for a reaction, which was introduced earlier in your study of chemistry. EXAMPLE 1 Calculate the free energy change for the complete combustion of one mole of methane, CH4(g), the main component of natural gas. Is this reaction spontaneous? SOLUTION We begin by writing the equation that represents this reaction. Recall that "complete combustion," or burning, is a reaction with oxygen from the atmosphere, forming carbon dioxide and water: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) Now we apply Equation 23.1 and then consult a table of thermodynamic values: ∆G° = (1 mol)[∆Gf° for CO2(g)] + (2 mol)[∆Gf° for H2O(l)] – (1 mol)[∆Gf° for CH4(g)] – (2 mol)[∆Gf° for O2(g)] = (1 mol) (–394.4 kJ/mol) + (2 mol) (–237.0 kJ/mol) – (1 mol) (–50.8 kJ/mol) – (2 mol) (0) = –817.6 kJ (Thermodynamic values are experimentally determined, so the values in your textbook may differ slightly from these.) The negative value of ∆G° indicates that the reaction is spontaneous. This matches our experiences in everyday life, where we have seen that natural gas burns spontaneously. Entropy The standard state entropy change for a reaction, ∆S°, can be calculated from data in thermodynamic tables in a manner similar to changes in enthalpy and free energy. ∆S° for a chemical reaction is ∆S° = ΣnS°products – ΣnS°reactants (Eq.2) Enthalpy, Entropy, and Free Energy Calculations A notable difference in ∆S° values is that we do not use "entropies of formation." This is a result of the Third Law of Thermodynamics, which defines a zero for entropy and thus allows us to calculate absolute entropy values. EXAMPLE 2 Determine the standard entropy change for the decomposition of one mole of solid calcium carbonate, forming solid calcium oxide and carbon dioxide gas. SOLUTION This is a straight-forward application of Equation 2, followed by substitution of the appropriate values from a table. CaCO3(s) → CaO(s) + CO2(g) ∆S° = (1 mol)[∆S° for CaO(s)] + (1 mol)[∆S° for CO2(g)] – (1 mol)[∆S° for CaCO3(s)] = (1 mol) (39.8 J/mol K) + (1 mol) (213.7 J/mol K) – (1 mol) (92.9 J/mol K) = 160.6 J/K Free Energy, Enthalpy, and Entropy By definition, Gibbs free energy is G ≡ H – TS (Eq.3) We are interested in the change of free energy associated with chemical reactions rather than absolute quantities, so we have ∆G = ∆H – ∆(TS) (Eq.4) and if we consider constant temperature processes, ∆G = ∆H – T∆S (Eq.5) and, finally, adding in standard state conditions, we have ∆G° = ∆H° – T∆S° (Eq.6) Equation 23.6, the Gibbs–Helmholtz equation, tells us that the standard free energy change depends on both the change in enthalpy and the change in entropy. We will explore this idea further during the workshop activities; for now, let's see how ∆G° can be calculated from ∆H° and ∆S° values. Enthalpy, Entropy, and Free Energy Calculations EXAMPLE 3 Calculate ∆G° for the reaction in Example 2, the decomposition of calcium carbonate, from ∆H° and ∆S° values. Enthalpy, Entropy, and Free Energy Calculations SOLUTION We have already calculated ∆S° for the reaction CaCO3(s) → CaO(s) + CO2(g) as 160.6 J/K. We can find ∆H° for the reaction in a similar manner: ∆H° = (1 mol)[∆H° for CaO(s)] + (1 mol)[∆H° for CO2(g)] – (1 mol)[∆H° for CaCO3(s)] = (1 mol) (–635.3 kJ/mol) + (1 mol) (–393.5 kJ/mol) – (1 mol) (–1207.0 kJ/mol) = 178.2 kJ Now we use Equation 23.6 to find the value of ∆G°: ∆G° = ∆H° – T∆S° = 178.2 kJ – 298.15 K × 160.6 J 1 kJ × = 130.3 kJ K 1000 J Notice that we used 298.15 K, or 25°C as the value of T. This is thermodynamic standard temperature. Also note how the ∆S° value was in joules per kelvin, while the ∆H° value was in kJ. The J ↔ kJ conversion must be accounted for in the calculation. Nonstandard Conditions and Free Energy Up to this point, we have only considered thermodynamic changes under standard conditions. ∆G at any conditions can be determined by ∆G = ∆G° + RT ln Q (Eq.7) where Q is the reaction quotient. This is the same quantity that was introduced while we were studying chemical equilibria. EXAMPLE 4 Consider the reaction of nitrogen monoxide and chlorine to form nitrosyl chloride: 2 NO(g) + Cl2(g) → 2 NOCl(g) a) Calculate ∆G° for the reaction. b) Calculate ∆G when pNO = 0.30 atm, pCl2 = 0.10 atm, and pNOCl = 0.45 atm. SOLUTION a) ∆G° is found by applying Equation 23.1: ∆G° = (2 mol)[∆Gf° for NOCl(g)] – (2 mol)[∆Gf° for NO(g)] – (1 mol)[∆Gf° for Cl2(g)] = (2 mol) (66.2 kJ/mol) – (2 mol) (86.6 kJ/mol) – (1 mol) (0) = –40.8 kJ Enthalpy, Entropy, and Free Energy Calculations b) ∆G at nonstandard conditions (the pressures are not 1 atm in this case) is found by applying Equation 23.7: ∆G = ∆G° + RT ln Q. Let's begin by calculating Q: Q = (pNOCl)2 (0.45)2 = = 23 (pNO)2 (pCl2 ) (0.30)2 (0.10) Now we can find ∆G: ∆G = ∆G° + RT ln Q = –40.8 kJ + 8.314 J 1 kJ × 298.15 K × ln 23 × = mol K 1000 J –33.0 kJ The Equilibrium Constant and Free Energy Starting from Equation Eq.7, ∆G = ∆G° + RT ln Q (Eq.7) and using the facts that at equilibrium, Q = K and ∆G = 0, 0 = ∆G° + RT ln K (Eq.8) and finally rearranging to isolate ∆G°, ∆G° = –RT ln K (Eq.9) we obtain the relationship between the change in free energy for a reaction and the equilibrium constant. EXAMPLE 5 Ksp for the reaction BaSO4(s) Ba2+(aq) + SO42–(aq) is 1.1 × 10–10. Use thermodynamic data to determine ∆G° for this reaction, and then calculate K from Equation 9. How do the K values compare? SOLUTION ∆G° is found in the usual manner: ∆G° = (1 mol)[∆Gf° for Ba2+(aq)] + (1 mol)[∆Gf° for SO42–(aq)] – (1 mol)[∆Gf° for BaSO4(s)] = (1 mol) (–560.8 kJ/mol) + (1 mol) (–744.5 kJ/mol) – (1 mol) (–1362.3 kJ/mol) = 57.0 kJ Enthalpy, Entropy, and Free Energy Calculations Now we can use Equation 23.9 to find K: –∆G° RT Let's get the coefficient of e first: ∆G° = –RT ln K ln K = K = e(–∆G°/RT) –∆G° K 1 1000 J = –57.0 kJ × × × = –23.0 RT 8.314 J 298.15 K kJ ↑ assume molar quantities K = e–23.0 = 1.0 × 10–10 The K calculated from ∆Gf° values agrees with the tabulated Ksp value to ±1 in the doubtful digit. Enthalpy, Entropy, and Free Energy Calculations Name: ______________________________ (To be done semi-independently with minimal consultation; include anyone’s name who helped you.) Thermochemistry Problems 1. Solid elemental sulfur can be produced, along with liquid water, by the reaction of hydrogen sulfide and sulfur dioxide gases. Calculate the standard free energy change for this reaction. 2. Nitrogen monoxide gas spontaneously decomposes into dinitrogen oxide and nitrogen dioxide gases. What is the standard entropy change for the decomposition of 3.0 mol of nitrogen monoxide? 3. A hypothetical reaction has ∆H° = –200.3 kJ and ∆S° = –77.0 J/K. Is this reaction spontaneous? Support your answer with the appropriate calculation. Enthalpy, Entropy, and Free Energy Calculations 4. Most of the direct energy needs of a cell are provided by the reaction of adenosine 5'-triphosphate (ATP) to form adenosine 5'-diphosphate (ADP) and hydrogen phosphate ion (Pi): ATP → ADP + Pi ∆G° = –30.0 kJ/mol for this reaction. What is ∆G when the concentrations in a cell are [ATP] = 3.2 × 10–3 M, [ADP] = 1.4 × 10–3 M, and [Pi] = 5.0 × 10–3 M? 5. Find the value of Ksp for iron(II) hydroxide from your textbook, then use Equation 9 to determine the value of ∆G° for the solution reaction of this slightly soluble solid. How does this value compare to the value determined by using ∆Gf° values? Enthalpy, Entropy, and Free Energy Calculations Name(s) ______________________________ (Turn-in one completed Worksheet per group with the names of each member who contributed.) Workshop: Enthalpy, Entropy, and Free Energy Calculations Questions 1–3, Suggested approach: select two partners, and have each member work on one of these three questions. When finished, discuss the results with the whole group and record your answers to turn in. 1. Consider the definition of Gibbs free energy, G ≡ H – TS, and the equations that can be derived from this definition. Construct a "truth table" showing all possible combinations of enthalpy and entropy changes for a chemical reaction and the resulting ability to predict the spontaneity of the reaction. Use your textbook to find an example of a reaction that fits each condition in your truth table. Enthalpy, Entropy, and Free Energy Calculations 2. Consider the definition of Gibbs free energy, G ≡ H – TS, and the equations that can be derived from this definition. Can a particular chemical reaction be nonspontaneous at one temperature but spontaneous at another temperature? If so, what criteria must be satisfied? If not, explain why. 3. Consider the equation that allows us to calculate free energy change under nonstandard conditions, ∆G = ∆G° + RT ln Q, and the equations that can be derived from this equation. Can the spontaneity of a reversible chemical reaction be determined solely from the equilibrium constant of that reaction? To answer this question, carefully consider the criteria for determining spontaneity and the relationship between the reaction quotient, Q, and the equilibrium constant, K. Explain your answer. Enthalpy, Entropy, and Free Energy Calculations Questions 4–8, Suggested approach: Complete each question as a round-robin. Use data from the thermodynamic tables in your textbook as necessary to answer the questions. 4. Determine the free energy change when 1.00 L of ethane, C2H6(g), at 25°C and 1.0 atm pressure, is completely oxidized. 5. Consider the decomposition of solid ammonium chloride to ammonia and hydrogen chloride gases. What do you predict for the sign of ∆S°? Consider a particulate-level explanation for your prediction. Calculate the value of ∆S° for the reaction, and compare it to your prediction. Now consider the decomposition of aqueous ammonium chloride to aqueous ammonia and hydrochloric acid. What do you predict for the sign of ∆S°? Why? Compare your prediction to the calculated value. 6. Consider these thermodynamic values for hypothetical compounds: Species (state) ∆Hf° (kJ/mol) A(g) –386.5 B(g) –139.9 Y(g) 33.6 Z(g) –295.2 S° (J/mol K) 177.0 234.8 277.1 301.3 Is the reaction A + B → Y + Z spontaneous? Is the reaction Y + Z → A + B spontaneous? How do you know? Enthalpy, Entropy, and Free Energy Calculations 7. One of the reaction steps for the metabolism of glucose in animals is not spontaneous: 2-phosphoglycerate → phosphoenolpyruvate ∆G°' = 1.7 kJ/mol (The prime indicates biochemistry standard state, which is the same as chemistry standard state except that biochemistry uses pH = 7.0 as a condition.) Will this reaction take place in a cell where [2-phosphoglycerate] = 2.3 × 10–4 M and [phosphoenolpyruvate] = 8.4 × 10–5 M? 8. Complete the blanks in the following statement. Significant quantities of both reactants and products are present at equilibrium for a reversible chemical reaction if ∆G° for that reaction is between _____ and _____.

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