Thermodynamics Lab: Calculating ΔG, ΔH and ΔS - PDF

Thermodynamics Lab: Calculating ΔG, ΔH and ΔS Purpose In this experiment you will use calorimetry to determine the enthalpy changes, ΔHrxn, for the dissolution of two chloride salts in water. You will

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Thermodynamics Lab: Calculating ΔG, ΔH and ΔS Purpose In this experiment you will use calorimetry to determine the enthalpy changes, ΔHrxn, for the dissolution of two chloride salts in water. You will then use textbook values to calculate ΔS o rxn and your experimental values to calculate ΔHrxn. Together, these will allow you to calculate the free-energy changes, ΔGrxn, for these two processes. Introduction The Gibbs-Helmholtz equation expresses the relationship between the free-energy change, the enthalpy change, and the entropy change at constant temperature and pressure: ΔG = ΔH - TΔS (equation 1) From knowing the value of ΔG, you may predict whether a process/reaction will be spontaneous at a certain temperature. A process is spontaneous if ΔG is negative (ΔG 0), nonspontaneous if ΔG is positive (ΔG 0), and at equilibrium if ΔG = 0. The enthalpy change, ΔHrxn, is the heat gained or lost by a system during a reaction carried out at constant pressure. Most reactions occur in several steps, with energy required (endothermic, positive ΔH) because energy is needed to break bonds, and energy released (exothermic, negative ΔH) because energy is released as new bonds are formed. ΔHrxn represents the total change in heat energy or enthalpy over the course of the reaction. In this experiment, you will use a coffee-cup calorimeter to determine the heat absorbed or released during the dissolution of ammonium chloride and the dissolution of calcium chloride. From observing the contents of the coffee-cup calorimeter, you will decide whether the dissolution processes are spontaneous or nonspontaneous. You will also calculate values of ΔG rxn to check your prediction. Discussion From the law of conservation of energy (energy is conserved) the total energy for the dissolution process is: q sys + q surr = 0 OR q sys = - q surr (equation 2) where qsys (or qrxn) represents the heat gained or lost by dissolving the solid, and qsurr (or qsoln) is the heat gained or lost by the solution in the calorimeter. Thus, heat energy is essentially transferred between the dissolved solid and the solution in the calorimeter. (For this experiment, we will consider the heat absorbed by the cup, probe, and surroundings to be negligible, so it is not included in the expression above.) The heat absorbed or released by the contents of the calorimeter is given by: q surr = (mass solution)*(specific heat of solution)*(δt) (equation 3) The mass of the solution is the sum of the masses of the water and salt placed in the calorimeter. (Recall that the density of water is 1.00 g/ml.) Because the solution is very dilute, the specific heat of the solution is basically equal to that of water, which is J/g C. To calculate change (Δ) for a variable it is always final minus initial. The heat of reaction, qsys, can then be calculated from combining equation 2 and equation 3. The molar enthalpy of reaction, ΔHrxn, will then be calculated by dividing the heat of reaction by the experimental number of moles of salt used in the experiment. ΔHrxn = q sys / moles salt (equation 4) You will need to calculate the ΔS o rxn values for the dissolution of solid ammonium chloride and calcium chloride using data from Appendix 2 in the back of your textbook. We do not have experimental data for this calculation, so we will use the textbook values and solve for ΔS o rxn like a homework problem. Finally you can calculate the experimental change in Gibb s Free Energy (ΔG) using the Gibbs-Helmholtz equation, equation 1, using the initial temperature for T, the experimental value for the enthalpy of reaction, and the textbook value for the entropy of reaction. Pre Lab Questions: In addition to title, purpose, and data/results tables, please include the following two calculations in your lab notebook and show your work for each. 1. Calculate the mass of ammonium chloride required to prepare 25 ml of a 2.0 M solution. 2. Calculate the mass of calcium chloride required to prepare 25 ml of a 2.0 M solution. 3. Define exothermic and endothermic: Exothermic: Endothermic: 4. When we look at the ΔH for a reaction, how can we identify (simply!) if the reaction is exo or endothermic? 5. The amount of heat, q, obtained in a reaction of 0.10 mole HCl(aq) with excess NaOH(aq) is 96 calories. (note: 1 calorie = Joules) a. Calculate the heat in terms of the number of Joules q = b. Calculate q in kj/mol HCl. q = c. If the reaction gives off the heat, is the reaction endo or exothermic? d. Therefore, what is ΔH for the reaction in kj/mol? kj/mol (Remember, use the correct sign for ΔH, that is consistent with your answer to the previous question!) 4. Answer the following questions: a. If the temperature of the water increases, did the water absorb heat or release heat? b. Therefore, will q H2O be +q or q? c. If the temperature of the water increases, did the reaction absorb heat of release heat? d. How are q H2O and q rxn related to one another? - They are exactly the same - They are numerically the same with opposite charges - They have nothing to do with one another e. Therefore, will the q rxn be +q or q? 5.Consider the combustion of acetylene gas: C2H2(g) + O2(g) CO2(g) + H2O(g) (unbalanced) a) Predict the signs of ΔH and ΔS b) Calculate ΔG for the combustion of one mole of acetylene by two different methods 6. Given the following thermochemical equations, using Hess law, calculate ΔH for the decomposition of one mole of acetylene (C 2 H 2 gas) into its elements in their stable/natural state at 25 C and 1 atm pressure.(hint: write the balanced chemical reaction for the decomposition of 1 mole of acetylene into its elements!) C 2 H 2 (g) + 5/2 O 2 (g) 2 CO 2 (g) + H 2 O (l) ΔH 1 = kj C (s) + O 2 (g) CO 2 (g) ΔH 2 = kj H 2 (g) + 1/2 O 2 (g) H 2 O (l) ΔH 3 = kj Balanced equation: ΔH rxn = kj Procedure 1. Use a 100-mL graduated cylinder to measure about 25 ml of deionized water and add to the clear plastic cup inside the Styrofoam cups. Record the exact volume of water used, paying attention to significant figures. 2. Tare out the weight of a plastic weighing cup. Remove the plastic cup from the balance and use a spatula to add the appropriate mass of ammonium chloride (calculated in pre-lab) to the weighing cup. The mass should be within ±0.2 grams of the calculated value. Record the exact mass used in your lab notebook. 3. Place a thermometer in the deionized water. Record the initial temperature of the water in degrees Celsius (Time = 0 sec). 4. Add the solid 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 three minutes. Record the temperature of the mixture every 10 seconds. Do not stir with the thermometer. Do not leave the thermometer in the apparatus unsupported it will be top-heavy and could fall over. 5. The highest (or lowest) temperature reached will be the final temperature T f. (Note: T f is NOT the temperature after 3 minutes, but the maximum (or minimum) temperature obtained during the three minutes.) 6. Observe the appearance of the salt solution in the calorimeter to see if the salt dissolved. Pour your salt solution in the sink. Rinse out your calorimeter, rinse the thermometer, then repeat the experiment using ammonium chloride again for trial number two. 7. Repeat all steps for calcium chloride for two trials. Clean-Up: CaCl 2 is hygroscopic and very corrosive to our balances. Please use the brush by the balance to clean up any spills immediately. Any spills left behind will result in points being deducted. Pour the salt solutions in the waste container. Rinse everything well with tap water followed by a quick DI water rinse. Return the measuring cups to the reagent stations. Clean your bench top. Put all equipment back exactly where you found it. Data and Results Record temperature data (every 10 seconds, for a maximum of 3 minutes) in your lab notebook. Two trials for each salt will be completed. The table on page 3 is an example of the data and results that you also need to record in your notebook. Please copy this table into your notebook twice (one for each salt). All data must be recorded in ink in your lab notebook as the reaction proceeds. Your calculations will also be completed in the lab notebook. You will conduct two trials for each salt. Pay attention to units, significant figures, and signs in your tables. Calculations Show all calculations for one trial for each salt. Example calculations should include: moles of salt mass of water mass of solution ΔT q sys ΔH rxn ΔS o rxn ΔG rxn In your conclusion in your notebook, summarize and discuss your calculated ΔH, ΔS o, and ΔG values for both salts. Also discuss at least 2 experimental sources of error and how they could have affected your results.
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