Thermochemistry: Enthalpy of Reaction Hess s Law

Transcription

1 Thermochemistry: Enthalpy of Reaction Hess s Law Objective Demonstrate Hess s Law for determining the enthalpy of formation for MgO by measuring temperature change for several reactions. Introduction The reaction between Mg and O 2 to form MgO is too dangerous for a basic chemistry laboratory. However, you can indirectly calculate the ΔH f of MgO by using Hess s Law. In this experiment, we will determine ΔH for reactions (1) and (2) below, and combine those vlaues with the given enthalpy for reaction (3) to determine the enthalpy of reaction (4). ΔH f for equation (3) is given and will be used in the calculation of ΔH f for MgO(s). Mg (s) + 2 H + (aq) Mg 2+ (aq) + H 2 (g) (1) MgO (s) + 2 H + (aq) Mg 2+ (aq) + H 2O (l) (2) H 2 (g) + ½ O 2 (g) H 2O (l) ΔH f = kj/mole (3) Mg (s) + ½ O 2 (g) MgO (s) (4) The reactions will be run at atmospheric pressure (we can assume that atmospheric pressure will be constant during this experiment). In general, the preferred units are kj for ΔH and kj per mole of compound for ΔH f. The sign of ΔH is relevant since a ΔH > 0 (positive) indicates heat absorbed (endothermic) and a negative ΔH indicates heat is released (exothermic). You will measure heat changes in this experiment by observing the temperature change of the reaction mixture. The reactants are mixed in a calorimeter (dual Styrofoam cup) and the change in the temperature of the reaction mixture is recorded until the temperature stabilizes (and all of the material is consumed). The heat released by the chemical reaction is all absorbed by the contents of the calorimeter, and the calorimeter equipment, resulting in an increase in the temperature of the vessel and its contents. The theoretical amount of heat released by chemical reaction in this experiment is equal to (ΔH) (n), where n is the number of moles of limiting reactant. Consistent with both thermodynamic convention and common usage, we define heat released (heat liberated) as a value less than zero (0) [and therefore equal to (ΔH) (n)].. The following example illustrates the definition. H 2 (g) + ½ O 2 (g) H 2O (l) ΔH f = kj/mole (5) If two moles of hydrogen react with one mole of oxygen to form 2 moles of water, the enthalpy (heat) change is calculated as follows: 1

2 ΔH = ΔH f(n) = kj/mole (2 moles) = kj The heat released is calculated as follows: Heat released = ΔH f (n) = ( kj/mole)(2 moles) = kj Thus we say the heat change of the system was minus kj because the system lost this amount of heat to the surroundings, but the heat released was kj. The heat released by the chemical reaction was absorbed by the experimental equipment (calorimeter plus contents). Part of the heat released by the reaction is gained by the reaction mixture (observed as a change in liquid temperature) and the rest is gained by the walls of the calorimeter (and the thermometer). The heat gained by the reaction mixture is equal to the product of [specific heat of the mixture (sp. ht.)] times [the mass of the reaction mixture] times [the change in temperature]. (Specific heat is defined as the amount of heat required to raise the temperature of one gram of substance one degree Celsius). ΔH rxn mixture = (mass) (sp. ht.) (ΔT) The heat gained by the calorimeter over small temperature ranges is a constant (Cc) times the change in temperature of the reaction mixture. The calorimeter constant, Cc, is specific for your calorimeter; therefore, it is important to use the same calorimeter for all the reactions. The units on the calorimeter constant are J/ C. If we have an adiabatic system, (no heat enters or escapes from the calorimeter) the heat change of the reaction can be approximated by the heat change of the reaction mixture and calorimeter. Mathematically, this is written: n ΔH = (mass) (sp. ht.) (ΔT) + (Cc) (ΔT) (6) (heat released by the reaction) = (heat gained by reaction mixture) + (heat gained by the calorimeter) To solve for ΔH for a reaction, all the other values in the above equation must be determined. The first objective will be to calibrate the calorimeter, and determine the calorimeter constant, Cc. This calibration is accomplished by producing a known quantity of heat (by a chemical reaction of H + and OH - ) from reaction (7) below and measuring ΔT. Since ΔH, the specific heat, and the quantities of NaOH and HCl are known (assuming that all volumes are additive), the only unknown in equation (6) is Cc, which you can now calculate. OH (aq) + H + (aq) H2O (l) ΔH = 57.7 kj/mole (7) Once the value of Cc is found, it can be used in reactions (1) and (2) to solve for the unknown heat of reaction, ΔH. 2

3 During this experiment you will measure and record the temperature at given time intervals. This data will be plotted on a graph of temperature versus time (as shown on the following page) to determine the maximum temperature for the reaction. Notice the temperature vs. time plot will be irregular until it stabilizes on a slowly decreasing temperature line. This irregular behavior is due to the calorimeter s inability to absorb heat as quickly as the reaction mixture releases it. However, accurate results can be obtained by extrapolating the temperature data to time zero (depicted in Figure 1). The correct temperature change, ΔT, is determined by subtracting the initial temperature (T i) from the extrapolated high temperature (T extrapolated). Skills: Quantitative Measurements Lab equipment: Scales, calorimeters Quantitative measurements Volumetric measurements Handling of chemicals Calculations Graphing of data Equipment & Reagents Calorimeter Distilled water Precision thermometer Clock or stopwatch 100 ml graduated cylinder Mg turnings 1M HCl MgO 1M NaOH Figure 1: Plot of Temperature versus time for the reaction of HCl with NaOH. 3

4 Procedure Part A: Calorimeter Calibration (finding Cc) 1. Choose a calorimeter and ensure the calorimeter and thermometer are clean and have been rinsed with distilled water. Using a 100 ml graduated cylinder, measure 50.0 ml of the approximately 1M HCl, then pour the HCl into the calorimeter. 2. Record the exact molarity found on the HCl storage bottle. 3. Insert the thermometer assembly and record the temperature when it stabilizes (less than 3 minutes). Measure all temperatures to the nearest 0.01 C with the aid of a special thermometer. WARNING: Precision thermometers are expensive. Treat them with care. DO NOT USE THE THERMOMETER TO STIR THE MIXTURE!!!!! If you break a thermometer, tell the instructor or lab assistant so they can have the mercury spill cleaned up. Also, HCl stains bench tops, books, and clothing; clean up spills immediately! 4. Rinse the graduated cylinder successively with tap water, distilled water, and 5 ml of 1M NaOH solution. 5. Measure 50.0 ml of NaOH solution in the graduated cylinder and, noting the time, pour the NaOH into the calorimeter. Immediately insert the thermometer assembly and gently swirl the reactants. 6. Record the temperature to the nearest 0.05 C at ½-minute intervals until a maximum is reached; then record at one minute intervals until enough data is obtained for an extrapolation (generally ~ 5 to 10 minutes). Gently swirl the reaction mixture before each reading. Heat of Reaction of Mg and HCl 1. Using a graduated cylinder, measure 50.0 ml of the 1M HCl and 50.0 ml of distilled water into the same clean, dry calorimeter used above and measure the stabilized temperature. 2. Weigh 0.20 to 0.30 g of Mg turnings to the nearest milligram on the balance. 3. Noting the time (t=0), drop the Mg into the calorimeter. Immediately put the thermometer assembly in place and swirl gently. DO NOT USE THE THERMOMETER TO STIR THE MIXTURE!!!!! 4. Record the temperature at ½-minute intervals until a maximum is reached. Then record at one minute intervals for at least 5 minutes. Gently swirl the reaction mixture between each reading. 4

5 Heat of Reaction of MgO and HCl 1. Repeat the above steps (1-4) using MgO in place of Mg. Use 0.50 to 0.60 g MgO, weighed to the nearest milligram. Be sure to use the same calorimeter. NOTE: You must swirl the calorimeter vigorously while reacting MgO and HCl since the MgO tends to form lumps at the bottom of the calorimeter and fails to dissolve. This could cause considerable error with appropriate impact on your grade. Hold the temperature assembly in the calorimeter so it will not break while swirling. When it appears that temperature changes have ceased, quickly look into the calorimeter. If there is still white solid at the bottom, the calorimeter must be swirled more vigorously until all the MgO has dissolved. Cleanup Clean your lab area, calorimeter, and glassware before being signed out by your lab assistant. Write-up As a critical part of the lab write-up, you will need to present your data from each of the experimental runs graphically (three separate plots), clearly indicating the final, initial, and change in temperature on each of the plots. To do this, we will rely on the Graphical Analysis computer program (or MS Excel, or any appropriate graphing program). MAKE SURE THAT YOU INCLUDE ALL APPROPRIATE INFORMATION ON THE PLOT TO ENSURE THAT THE PLOTS STAND ALONE AND CONVEY APPROPRIATE INFORMATION. If I presented your plots back to you at a later date, could you understand the data/information presented? Each student must include plots of their data as a part of the lab write-up. Calculations You must show the appropriate representative calculations (using your data) for each of the calculations below. These calculations may be NEATLY hand-written, and forego the routine type-written requirement. It goes without saying that you must explicitly indicate the units used for each calculation. 1. Determine the calorimeter constant, Cc, for the calorimeter you used. 2. Determine ΔH for the reaction of Mg(s) with HCl(aq). (Equation 1) 3. Determine ΔH for the reaction of MgO(s) with HCl(aq). (Equation 2) 4. Calculate ΔH f for MgO(s). (Equation 4) 5

6 Post Lab Questions (Answer the following questions, making sure to give me a good indication of the question, your answer, AND your rationale). 1. If heat is conserved, why does the calorimeter cool down? 2. If ΔH were not a state function, would this technique be valid? Explain. 3. How would the overall (net) results be affected if we doubled the amount of MgO or Mg? 4. If all of the MgO in the calorimeter has not reacted when you stop taking data, how will your final results be affected? 6

7 5. Report: Hess's Law Name Lab Partner(s) Section Date performed Part I: Calorimeter Calibration Data Volume of acid ml Temperature of acid and calorimeter C Volume of base ml Temperature of base C Avg. Temp of acid and base (before mixing) C Temperature vs. time after mixing of HCl and NaOH: (The first time/temp block should be the time = 0 (average temperature from above). Time Temp Time Temp Time Temp Time Temp Time Temp 0 Specific heat of reaction mixture J/g C (Given) 7

8 Part II & III: Heats of Reaction Data Part II Part III Mg & HCl MgO & HCl Volume of acid in calorimeter ml ml Volume of water in calorimeter ml ml Temperature of solution at equilibrium (Ti) C C Mass of solid reactant g g Temperature vs. time after mixing: Mg & HCl Time Temp Time Temp Time Temp Time Temp Time Temp 0 MgO & HCl Time Temp Time Temp Time Temp 0 Specific heat of reaction mixture Mg & HCl J/g C (Given) MgO & HCl J/g C (Given) NOTE: The sign of ΔH must be taken into account for all calculations. 8