|
Sarah C. 04/13/09 Experiment 13.1 Determining the ΔH of a Chemical
Reaction A. Purpose: Purpose: The purpose of this experiment is to demonstrate to the experimenter the steps taken by scientists to conclude the value of the ΔH of a particular reaction. Over the history of thermodynamics, several terms have been used to denote what is now known as the enthalpy of a system. Originally, it was thought that the word "enthalpy" was created by Benoit Paul Émile Clapeyron and Rudolf Clausius through the publishing of the Clausius-Clapeyron relation in The Mollier Steam Tables and Diagrams in 1827. Josiah Willard Gibbs introduced a "heat function for constant pressure" in 1875,[2] although the word enthalpy does not appear in any of Gibbs's works.[citation needed] In 1909, Keith Landler discussed Gibbs's work on this "heat function" and noted that Heike Kamerlingh Onnes had coined the modern name from the Greek word "enthalpos" (ενθαλπος) meaning "to put heat into." [note 1] Original definition: The thermodynamic potential H was introduced by the Dutch physicist Kamerlingh Onnes in early 20th century in the following form: H = E + pV Where E represents the energy of the system. In the absence of an external field, the enthalpy may be defined, as it is generally known, by: H = U + pV Where (all units given in SI) H is the enthalpy (in joules), U is the internal energy (in joules), p is the pressure of the system, (in pascals), and V is the volume, (in cubic meters). The form pV
(sometimes called "flow work") is motivated by the following
example of an isobaric process. Gas undergoing combustion in a cylinder
pushes a piston, maintaining constant pressure p and adding heat to the gas.
The force is calculated from the area A of the piston and definition of
pressure p = F/A: the force is F = This experiment hopes to show the steps involved when attempting to find the ΔH, or change in enthalpy, of any chemical reaction and that these steps can also be performed by a mere student; these procedures are not reserved for experts in their practice. This topic, videlicet, changes in enthalpy, is of interest to science because every substance in nature contains stored energy. To be able to measure this enthalpy, or the change of which, is necessary to understanding energy transfers in chemical equations. Hypothesis: If the experimenter places the vinegar into the makeshift calorimeter and measures the temperature with her thermometer, then the temperature will read lower than room temperature. B. Equipment: 1. Two Styrofoam coffee cups 2. Thermometer 3. Vinegar 4. Mass scale 5. Measuring tablespoon and ˝ teaspoon 6. Lye 7. Safety goggles C. Procedure: 1. In order to perform this experiment, a small amount of lye must be measured. The problem is, the mass scale cannot measure small masses very accurately. Thus, the experimenter must find how many grams are in a teaspoon of lye; then, using measuring spoons, find the mass of a small amount of lye. To do this, measure the mass of ten tablespoons of lye. 2. Remembering that each tablespoon represents three tablespoons, realize that the mass just measured is the mass of 30 tablespoons of lye. Now take that mass and divide it by 30. The result is the mass of one teaspoon of lye. 3. Calculate the mass of ˝ teaspoon of lye by dividing the mass of one teaspoon by two. 4. Nest one coffee cup inside the other to make a calorimeter. 5. Pour 100.0 mL of vinegar into the calorimeter. 6. Place the thermometer in the vinegar and let it sit for three minutes. After the time has elapsed, record the temperature of the vinegar. This is the initial temperature of the experiment. 7. Now add ˝ teaspoon of lye to the vinegar and begin stirring the mixture with the thermometer. 8. This reaction is exothermic; thus the temperature should begin to rise. Read the temperature every 30 seconds. Once the temperature falls or stays constant for two consecutive readings, end the experiment. 9. Now, begin the calculations. The equation is q = m•c•ΔT. Ignore the calorimeter. 10. The specific heat of vinegar is 4.1 J/g*C. 11. The mass of the vinegar in the calorimeter is its volume times its density. In addition to this number, add the mass of ˝ teaspoon of lye. 12. Subtract the initial temperature in the experiment from the final temperature to get ΔT. 13. Now calculate the heat absorbed by the calorimeter. Take the specific heat given in step 11 times the mass calculated in step 12, then multiply it by the ΔT calculated in step 13. This is the heat absorbed by the calorimeter’s contents. 14. Since the heat absorbed by the calorimeter’s contents came from the chemical reaction, this number is the ΔH of the reaction. 15. This ΔH is the amount of energy released when the mass of lye that was used in the experiment reacted with vinegar. Now determine the ΔH per mole of NaOH by taking the ΔH and divide it by number of moles of NaOH used in the experiment. 16. Take the mass of lye and divide it by the molar mass of NaOH. That is the number of moles of NaOH that were used in the reaction. 17. Divide the ΔH calculated in step 14 by the number of moles calculated in step 17. The result is the ΔH of the reaction in J/mole. 18. Clean up the mess. D. Observations: 1. 10 Tbs lye = 212 g 2. 1 tsp lye = 7.07 g 3. ˝ tsp lye = 3.53 g 4. Starting temperature = 18.3 *C 5. Ending temperature = 32.9 *C 6. Mass of vinegar = 99g 7. Mass of calorimeter = 103 g 8. ΔT = 14.6 *C 9. q = m•c•ΔT 10. qcalorimeter = 6200 J 11. ΔH = 6200 J 12. 23.0 + 1.01 + 16.0 = 40.0 amu 13. 3.53 g NaOH • 1 mol NaOH/40.0 g NaOH = .0883 mol NaOH 14. 6200 J/.0883 mol = 7.0 x 104 J/mol E. Conclusions: The above hypothesis was supported with the data gathered: when the experimenter placed the vinegar into the makeshift calorimeter and measured the temperature with her thermometer, the temperature read lower than room temperature. The thermometer read 18.3 *C, room temperature being between 20 and 25*C. The experimenter expected this result as, when her hand came in contact with the vinegar, it felt quite cold to the touch. Enthalpy [is] defined as the sum of the internal energy and the product of pressure and volume. In other words, H = U + PV. The justification for defining this variable is really only a matter of convenience, because we often find that the sum U + PV occurs in thermodynamic equations. This isn’t surprising, because the work done by a quantity of gas depends on the product of pressure times volume. When a gas expands quasi-statically at constant pressure, the incremental work δW done on the boundary is PdV, so from the energy equation dU = δQ – δW we have δQ = dU + PdV. Noting that, at constant pressure, dH = dU + PdV, it follows that δQ = dH for this process. This explains why enthalpy is often a convenient state variable, especially in open systems. Obviously enthalpy has units of energy, but it doesn’t necessarily have a direct physical interpretation as a quantity of heat. In other words, enthalpy is not any specific form of energy, it is just a defined variable that often simplifies the calculations in the solution of practical thermodynamic problems (MathPages). This experiment could have been improved by using a more precise measuring device for the lye. The inability to precisely measure the amount of lye used alters the answer, however subtly. However, since the experimenter is not in a professional situation and performed this experiment merely to gain a better understanding of its inner workings, this deficiency is of little or no consequence. Ideas for further research were generated by the possibility of procuring a more precise measuring unit. The experimenter wonders where such a unit would be found; a puzzle, as conventional mass scales are not precise enough to measure such an amount. F. Bibliography: "Energy, Entropy, Enthalpy." MathPages.com. MathPages. 03/28/09. Domain: http://www.mathpages.com/ Document: home/kmath184/kmath184.htm "Enthalpy." Wikipedia, The Free Encyclopedia. 16 Aug 2009, 17:47 UTC. 03/16/09 Domain: http://en.wikipedia.org/ Document: wiki/Enthalpy#History Rosenoff, Steven. Class Lecture. 03/24/09 Wile, Dr. Jay L. Exploring Creation with Chemistry, 2nd ed. CJK: Apologia Educational Ministries, 2007. |