Thermodynamics

Thermodynamics Introduction: Gibbs Free Energy, G, can be used to determine if a reaction is spontaneous or not. A negative value of G indicates that a given reaction is spontaneous at the measured conditions and will proceed in the forward direction. G can be calculated using the Gibbs-Helmholtz equation: G = H – TS. In this experiment, H (enthalpy) will be calculated from the temperature change of the reaction. S (entropy) will be calculated using standard entropy values from the textbook. The reactions used in this lab will be the dissolution (dissociation) of two salts in water. It will be important to differentiate between the system (or reaction) and the surroundings in this experiment. The calorimeter used in this experiment is assumed to be a closed, isolated container that does not lose any heat to the environment. Therefore, all heat exchanges are assumed to take place between the system and the surroundings. The dissolution of each salt is the system; water is the surroundings. It is the temperature change of the water (the surroundings) that will be measured over time. Refer to Sections 5.2, 5.3, 16.1, 16.2, and 16.4 of Openstax Chemistry for information on calorimetry and calculating q, enthalpy, spontaneity, entropy, and free energy. Equations to use for the calculations: Heat of solution: qsoln = mass x specific heat x T, where mass = mass of solution specific heat = 4.184 J/g oC T = Tfinal - Tinitial Tfinal is the maximum or minimum temperature reached Tinitial is the initial temperature of water before adding salt Law of Conservation of Energy: qrxn + qsoln = 0 → qrxn = - qsoln Density of water: 1.00 g/mL; Use to find mass of measured volume Enthalpy of reaction: DH rxn = q rxn moles of salt Entropy of reaction: Sorxn = n So (products) – n So (reactants) Free Energy of reaction: Grxn = Hrxn - TSrxn (where T = Tinitial in Kelvin) Materials: Styrofoam cup calorimeter Digital thermometer 100 mL graduated cylinder plastic weighing cup timer DI H2O solid calcium chloride, CaCl2 solid ammonium chloride, NH4Cl page 1 of 9 Procedure: Note: Thoroughly clean up any spilled salts. CaCl2 will corrode the metal on the balances. Use the brush to sweep off the metal surface of the balances immediately. Any spills left behind might result in points being deducted! Do NOT return excess reagent to the salt container. It must go into the designated waste. 1. Place the nested Styrofoam cups into a 400-mL beaker. Use a 100-mL graduated cylinder to measure about 25 mL of deionized water and add to the blue plastic cup inside the Styrofoam cups. Record the exact volume of water used. 2. Tare out (zero) the mass of the weighing cup. Remove the weighing cup from the balance and use a spatula to add the appropriate mass of ammonium chloride (calculated in pre-lab). The mass should be within 0.20 grams of the calculated value. NOTE: DO NOT ADD SALT TO THE WEIGHING CUP WHILE IT IS ON THE BALANCE! Refer to the “Using an analytical balance”, “Transferring solids”, and “Preparing solutions” techniques. 3. Place a thermometer in the DI water. This will be your initial temperature. 4. Add the salt to the water in the calorimeter and replace the lid. Stir the solution vigorously by swirling the beaker and contents, carefully holding the lid and thermometer in place, for two minutes. Record the temperature of the mixture every 10 seconds. 5. The highest (or lowest) temperature reached will be the final temperature, Tf. Note: Tf is NOT the temperature after 2 minutes, but the maximum (or minimum) temperature obtained. 6. Be sure to record your observations of the appearance of the salt solution in the calorimeter at the end of two minutes. Pour your salt solution in the waste container in the hood. Rinse, dry, and reuse the plastic weighing cup and the plastic coffee cup for your second trial. 7. Repeat all steps for a second trial of ammonium chloride. Repeat all steps for two trials of calcium chloride. Waste Disposal: Pour the salt solutions in the waste container. Clean-Up: Rinse everything well with tap water followed by a quick DI water rinse. Wipe your benchtop with a damp paper towel. Put all equipment back where you found it. page 2 of 9 Thermodynamics Pre-Lab Questions and Calculations You will complete this quiz in Canvas 1 hour before your lab period. This page will not be turned in or graded. You may use this page to set up your calculations before you take the Canvas quiz. 1. Calculate the mass of ammonium chloride required to prepare 25 mL of a 2.0 M solution. Show your work. Record this value on your data table (page 4) to reference during lab. 2. Calculate the mass of calcium chloride required to prepare 25 mL of a 2.0 M solution. Show your work. Record this value on your data table (page 4) to reference during lab. 3. Write the balanced equation for the dissolution of ammonium chloride in water. 4. Write the balanced equation for the dissolution of calcium chloride in water. 5. What does a negative value of H indicate about a reaction? 6. What does a negative value of S indicate about a reaction? 7. What does a negative value of G indicate about a reaction? 8. Calculate the heat of solution (qsoln) if the mass of solution is 45.6 g, specific heat is 3.98 J/g oC, and T is -34.5oC. 9. What is the heat of reaction (qrxn) for the above question? 10. If 0.612 moles of salt were used in the reaction, what is Hrxn? page 3 of 9 Name: _____________________________ Partners: ________________________ Thermodynamics Lab Report Turn in pages 4-7 as your Lab Report. Data and Calculations Table 1: Mass, volume, and temperatures Ammonium Chloride Calcium Chloride Calculated mass from prelab, g Calculated mass from prelab, g Mass of salt, g Mass of salt, g Volume of water, mL Volume of water, mL Initial temp. of H2O,Ti (C) Initial temp. of H2O,Ti (C) Time Temperature, oC, Trial 1 Temperature, oC, Trial 2 Time 0:00 0:00 0:10 0:10 0:20 0:20 0:30 0:30 0:40 0:40 0:50 0:50 1:00 1:00 1:10 1:10 1:20 1:20 1:30 1:30 1:40 1:40 1:50 1:50 2:00 2:00 Temperature, oC, Trial 1 Temperature, oC, Trial 2 Table 2: Observations of each salt after reaction and dissolution reaction in water: Observation after reaction - ammonium chloride: Observation after reaction – calcium chloride: Balanced equation for dissolution of ammonium chloride: Balanced equation for dissolution of calcium chloride: Instructor Initials: _________ page 4 of 9 Calculations for NH4Cl: Show a sample calculation for one trial for a) moles of salt, b) mass of solution, c) T, d) qrxn, e) H, f) So (Use values in Table 3 on page 6 for this calculation.), and g) G. Refer to Results Table 4 for units of each quantity. Instructor Initials: _________ Calculations for CaCl2: Show a sample calculation for one trial for a) moles of salt, b) mass of solution, c) T, d) qrxn, e) H, f) So (Use values in Table 3 on page 6 for this calculation.), and g) G. Refer to Results Table 4 for units of each quantity. Instructor Initials: _________ page 5 of 9 Table 3: Thermodynamic data from textbook. Hfo (kJ/mol) Gfo (kJ/mol) NH4Cl (s) -314.4 -202.8 NH4+ (aq) -132.5 -79.3 Cl (aq) -167.2 -131.2 CaCl2 (s) -795.8 -748.1 2+ Ca (aq) -542.9 -553.0 So (J/K∙mol) 94.6 113.4 56.5 104.6 -55.2 Results: Table 4: Results of salt dissolution calculations Ammonium Chloride Trial 1 Trial 2 Calcium Chloride Trial 1 Trial 2 Moles of salt (mol) Mass of Water (g) Mass of Solution (g) Final temperature,Tf (C) ΔT (C) qrxn (J) ΔHrxn (kJ/mol) ΔSorxn (J/mol K) calculated from table 3 ΔGrxn (kJ/mol) Use values from Table 3 (top of this page) to calculate standard values of Horxn and Gorxn for NH4Cl. Hrxn: Grxn: page 6 of 9 How do your calculated ΔHrxn and ΔGrxn values using experimental data compare to the standard values? Using your results for trial 1, calculate the temperature (in Kelvin) CaCl2 becomes spontaneous. Is the dissolution of CaCl2 spontaneous above or below the temperature calculated above? Was this process spontaneous at room temperature? Conclusion: In the space below, summarize your average calculated values for H, S, and G for each salt. Describe what the sign of each value tells you about the dissolution of that salt. page 7 of 9 Post-Lab Questions – These questions will not be graded as part of your lab report grade. You will be responsible for the information in these questions and able to answer these or similar questions on the post-lab quiz at the start of next week’s lab period. Dissolution of ammonium chloride: 1. In the experiment, identify the system ______________ and the surroundings _______________. 2. Which one gains heat in this experiment? _______________ 3. Is the system endothermic or exothermic? ___________________ 4. Explain how the observed temperature change verifies your answer to #3. 5. From the temperature change obtained for the solution in the calorimeter, what must be the sign for Hrxn? _____________________ 6. a) Based on your observations at the end of the reaction, is the dissolution of ammonium chloride spontaneous or non-spontaneous at room temperature? ___________________________ b) Based on your observations at the end of the reaction, is G > 0 or < 0 for this process at room temperature? _______ 7. Based on your calculated change in entropy, does the dissolution of ammonium chloride create more order or disorder? __________________ 8. Is this salt soluble at all temperatures? Explain based on the signs of the enthalpy change and entropy change for the dissolution of ammonium chloride. 9. Calculate the temperature above or below which the salt will not dissolve, if applicable. Indicate if the solid will dissolve above or below this temperature. page 8 of 9 Dissolution of calcium chloride: 1. In the experiment, identify the system ______________ and the surroundings _______________. 2. Which one gains heat in this experiment? _______________ 3. From the temperature change obtained for the solution in the calorimeter, what must be the sign for Hrxn? ______________ 4. a) Based on your observations at the end of the reaction, is the dissolution of calcium chloride spontaneous or non-spontaneous at room temperature? ___________________________ b) Based on your observations at the end of the reaction, is G > 0 or < 0 for this process at room temperature? _______ 5. Based on your calculated change in entropy, does the dissolution of calcium chloride create more order or disorder? __________________ 6. Is this salt soluble at all temperatures? Explain based on the signs of the enthalpy change and entropy change for the dissolution of calcium chloride. 7. Calculate the temperature above or below which the salt will not dissolve, if applicable. Indicate if the solid will dissolve above or below this temperature. Discuss at least 2 sources of error. How did they affect your results and how would you correct them if you were to repeat the experiment? page 9 of 9