Determining Heat and Enthalpy of Reactions: A Calorimetry Experiment, Study notes of Law

Download Determining Heat and Enthalpy of Reactions: A Calorimetry Experiment and more Study notes Law in PDF only on Docsity! 93  Enthalpy of Reaction  E x p e r i m e n t 8 Lecture and Lab Skills Emphasized • Calculating the heat and enthalpy of reactions. • Writing net ionic equations. • Using Hess’s law to determine the enthalpy of a reaction. • Defining calorimetry. In the Lab • Students will work in pairs. • Record your procedure and original data in your lab notebook along with your calculations. • Report data collected and subsequent calculations to www.chem21labs.com. • All equipment should be returned to the correct location after use. Waste • All solutions should be should be disposed of in the acid–base waste container. • Styrofoam cups should be rinsed, reused, and returned at the end of the lab session. Do not throw away.  94  E x p e r i m e n t 8 • Enthalpy of Reaction Expt. 8 Safety • Concentrated acids and bases can cause serious burns. Use with caution. • Gloves and safety goggles are mandatory when anyone is performing an experiment in the lab. • Long pants, closed-toed shoes, and shirts with sleeves. Clothing is expected to reduce the exposure of bare skin to potential chemical splashes. • Always wash your hands before leaving the laboratory. Additional information can be found at http://genchemlab.wordpress.com/8-enthalpy/. To continue to show your analytical ability, your com- pany has asked you to determine the enthalpy for a chemical reaction using three different methods. Re- member, you will need to identify these three methods and be careful of your laboratory technique to ensure your results come out the same. All chemical reactions involve an exchange of heat en- ergy; therefore, it is tempting to plan to follow a reaction by measuring the enthalpy change (H). However, it is often not possible to directly measure the heat energy change of the reactants and products (the system). We can measure the heat change that occurs in the sur- roundings by monitoring temperature changes. A calo- rimeter is a device used to measure the heat changes when two or more substances are combined, whether it be two substances at different temperatures or two substances that will undergo a chemical reaction. The key feature of a calorimeter is that it is insulated enough to minimize the loss of heat to the surroundings. A Styrofoam cup is an effective calorimeter because it acts as a good insulator. Inserting the calorimeter into a beaker provides extra insulation and stability to the cup. Ideally, a calorimeter would not absorb any heat, but in reality it does. By performing a calibration of the calorimeter, we will be able to determine the amount of heat absorbed by the calorimeter and will account for that loss when determining the heat change for our chemical reactions. In this experiment, we will first determine the heat capacity for our calorimeter by mixing samples of hot and cold water together. Based on the difference in the temperature of the two water samples, we can determine the amount of heat lost to the calorimeter. This will allow us to account for the amount of heat lost to the calorimeter when we are measuring the heat change for three acid–base reactions. If we assume that no heat is lost to the surroundings, then the following expression is true q cold water  qhot water  qcal  0 (1) and this can be rearranged to the following, qcal  qhw  qcw (2) which will allow us to determine the heat of the calo- rimeter. There are two formulas which we will use to calculate the individual terms in the equation. The first formula, which is shown in equation 3, is used when you have varying masses of a substance but you know the specific heat of the substance. The specific heat, s, has been determined for many substances (i.e., water, iron, etc.) and can be found in a variety of reference sources. q  msT (3) The second formula, shown in equation 4, is used for determining the heat when you are dealing with an object with a fixed mass whose composition will not change, such as a calorimeter. The value of the heat capacity, C, is unique to a particular item and must be determined experimentally for that object. q  CT (4) 97  E x p e r i m e n t 8 • Enthalpy of Reaction  Expt. 8 9. Pour approximately 25 mL of water into the cup. 10. Determine the exact mass of the cup and water. 11. Place the cup with the water inside the 250 mL bea- ker and lay the cardboard square on top as shown in Figure 8.2. 12. Determine the mass of your 100 mL beaker. 13. Add approximately 25 mL of water into the beaker and then determine the exact mass of the beaker and water. 14. Using the hot plate, gently heat the water in the bea- ker to approximately 50–60°C. Use the thermometer to measure the temperature of the hot water. Steps 15–18 should be done quickly. 15. Once the temperature of the hot water is constant, record the initial temperature of the hot water in your lab notebook. 16. Insert the temperature probe into the clamp, through the hole in the cardboard square, and into the calorimeter (see Figure 8.2). Be careful not to poke a hole in the bottom of the styrofoam cup. 17. Press “START/STOP” on the workstation. 18. Wait until the temperature in the calorimeter is constant for a few seconds. Without removing the temperature probe from the water, pour the hot water into the styrofoam cup with the cold water. Replace the cardboard lid. Gently swirl the calorimeter to mix the solution, being careful not to lose any of the water. 19. Continue collecting data until the temperature starts to drop. The maximum temperature will be the final temperature. Press “START/STOP” to stop collecting data. 20. Press “FILE OPTIONS” on the workstation and select “SAVE DATA.” When prompted for a file name, use “010.” You must use the correct file name so that the data is associated with your Chem21 account. If you save your file with the wrong name, repeat the save with the correct name. 21. Check with your TA to make sure the file was saved and uploaded correctly. 22. Pour the water in the drain and dry the styrofoam cup. Record the exact volumes of all liquids used. Estimate one decimal place beyond the markings on the glassware. Part I—NaOH and HCl 23. Using a clean, dry 50 mL graduated cylinder, mea- sure approximately 25 mL of HCl and record the exact volume used. 24. Pour the HCl into the cup. 25. Place the cup with the HCl inside the 250 mL beaker and lay the cardboard square on top as shown in Figure 8.2. 26. Insert the temperature probe into the clamp, through the hole in the cardboard square, and into the calorimeter (see Figure 8.2). Be careful not to poke a hole in the bottom of the styrofoam cup. 27. Using a clean, dry 50 mL graduated cylinder, mea- sure approximately 25 mL of NaOH and record the EXACT VOLUME used. 28. Press “START/STOP” on the workstation. 29. Wait 5–10 seconds and then quickly, but carefully, pour the NaOH solution into the calorimeter with the HCl solution. 30. Gently swirl the calorimeter to mix the solution, being careful not to lose any of the solution. 31. The temperature will rise, plateau, and then drop. Once it is clear that the temperature is dropping, press “START/STOP” on the workstation. 32. T will be determined based on the initial tempera- ture of the solution and the maximum temperature. Since the solutions and calorimeter are all at room temperature, we can assume that the initial and final temperatures for all three are the same.  98  E x p e r i m e n t 8 • Enthalpy of Reaction Expt. 8 33. Press “FILE OPTIONS” on the workstation and select “SAVE DATA.” When prompted for a file name, use “001.” You must use the correct file name so that the data is associated with your Chem21 ac- count. If you save your file with the wrong name, repeat the save with the correct name. If you use a previous file name from this experiment, your data will be overwritten. 34. Check with your TA to make sure the file was saved and uploaded correctly. 35. Dispose of the solution and clean the cup for the next use. Part II—NaOH and NH4Cl 36. Repeat steps 24–35 with NaOH and NH4Cl solu- tions. Put the NH4Cl in the cup first and add the NaOH. 37. Be aware that this reaction has a smaller T than Part I. 38. Save the file as “002.” If you use a previous file name from this experiment, your data will be overwritten. Part III—HCl and NH3 39. Repeat steps 24–35 with HCl and NH3 solutions. 40. Save the file as “003.” If you use a previous file name from this experiment, your data will be overwritten. Data Analysis 1. What was the mass and value of T for the calo- rimeter? 2. What was the mass and value of T for the cold water? 3. What was the mass and value of T for the hot water? 4. Determine the heat change (q) for the cold and hot water. 5. Calculate the heat change for the calorimeter (pay attention to the sign on the value of q). 6. Using T for the calorimeter, determine the heat capacity of the calorimeter. 7. Calculate the heat change for each of the three reactions. Use 1.03 g/mL for the density of all solutions. 8. For each reactant, determine the limiting reagent (all reactions have a 1:1 molar ratio) and use the amount of that reactant, in moles, to determine the enthalpy change for each of the three reactions. 9. Use the experimental data for reactions 1 and 2 along with Hess’s law to determine the enthalpy of Reaction III. 10. Using the two enthalpy values calculated previously (questions 8 and 9), determine the percent error in each of these values compared to the accepted value calculated in the pre-lab exercise.