CHEM 131L Principles of Chemistry Laboratory Lab 10: The Enthalpy of Reaction an

CHEM 131L Principles of Chemistry Laboratory Lab 10: The Enthalpy of Reaction an

CHEM 131L Principles of Chemistry Laboratory Lab 10: The Enthalpy of Reaction and Hess’s Law Remote Instruction Version Introduction The enthalpy of a system, H, represents its total heat energy. In a chemical reaction, when the reactants have a greater enthalpy than the products, the excess heat energy is released to the surroundings, causing an increase in temperature. This is called an exothermic reaction. Conversely, in an endothermic reaction the products have a greater enthalpy than the reactants. In these reactions, the extra energy of the products must be absorbed from the surroundings, causing a decrease in temperature. In either case, the enthalpy change of the chemical reaction is simply the difference between the enthalpies of the products and the reactants: ?H = Hproducts – Hreactants With constant-pressure calorimetry, the enthalpy change of a chemical reaction can be determined by measuring the amount of heat, q, transferred to the calorimeter, and dividing by the number of moles, n, of the limiting reactant: ?H = qrxn/nlim The change in enthalpy of a reaction is expressed in units of kJ/mol. Since enthalpy is a state function, the change in enthalpy in going from some initial state to some final state is independent of the pathway. This means that in going from a particular set of reactants to products, the ?H is the same whether the reaction takes place in one step or in a series of steps. This principle is known as Hess’s law and can be illustrated by examining the reaction between nitrogen and oxygen gas to form nitrogen dioxide. This reaction can be written in one step, where the enthalpy change is represented by ?H1: N2 (g) + 2 O2 (g) ? 2 NO2 (g) ?H1 = 68 kJ However, the reaction can also be carried out in two distinct steps, with the enthalpy changes designated by ?H2 and ?H3: N2 (g) + O2 (g) ? 2 NO (g) ?H2 = 180 kJ 2 NO (g) + O2 (g) ? 2 NO2 (g) ?H3 = -112 kJ Net reaction: N2 (g) + 2 O2 (g) ? 2 NO2 (g) ?H2 + ?H3 = 68 kJ Note that the sum of the two steps gives the net, or overall, reaction and that ?H1 = ?H2 + ?H3 = 68 kJ This experiment will focus on the enthalpy change associated with three reactions involving NaOH, and the verification of Hess’s law. Procedure You will be using virtual simulations to perform this experiment. Please go to the website: https://media.pearsoncmg.com/bc/bc_0media_chem/chem_sim/calorimetry/Calor.php You will see the screen below. Click on the ‘experiment’ tab. When the experiment screen appears, you will see the following box: Click on ‘Run demonstration’ and follow the steps. The step-by-step description for running the demonstration will appear in the panel on the left side of your screen. Please proceed with each step as shown in the panel. The demonstration will show you how to mix two solutions or a solid and a solution, measuring temperature change for every reaction and plotting the graph. Follow the steps in this panel Once you have finished running the demonstration (it will take about 8-10 minutes), you are ready to run the experiment. Use the data tables later in the procedure to make note of your observations and data. Part 1 – The Molar Heat of Solution of Sodium Hydroxide 1. In the first column, select solids. Then, in the drop-down menu, select NaOH. 2. Choose any mass of NaOH, but keep the initial temperature at 20 °C. 3. Record the mass and temperature of NaOH exactly as shown in the boxes. 4. Click on the ‘next’ button in the panel on the left-hand side to proceed to the next column. 5. In the next column, select water under the ‘liquids’ tab. 6. Choose your initial mass of water, but keep the initial temperature at 20 °C. 7. Record the mass and temperature of water exactly as shown in the boxes. 8. Under ‘Run experiment’, select the options for ‘show graph view’ and ‘show microscopic view’. Hit the start button. 9. After the simulation is finished running, a temperature will be displayed on top of the calorimeter. Make note of this temperature exactly as shown. 4 Record the initial and final temperatures for the trial. You may approximate the specific heat of the solution as that of pure water. Once you have noted all the numbers, click on the ‘Reset’ button on the bottom right corner to run the next experiment. Part 1 – The Molar Heat of Solution of Sodium Hydroxide Specific heat of solution Mass of NaOH pellets (grams) Initial water temperature (°C) Final solution temperature (°C) Temperature change (?°C) Mass of solution (in grams) Part 2 – The Molar Heat of Neutralization 1. In the first column, click on the ‘solutions’ tab. In the drop-down menu, select NaOH (aq). Enter any volume and molarity you prefer. Record the volume and molarity of NaOH exactly as shown in the boxes. Click on the ‘next’ button in the panel on the left-hand side to proceed to the next column. 2. In the second column, click on the ‘solutions’ tab. Then, in the drop-down menu, select HCl (aq), and again choose your volume and molarity. Record the volume and molarity of HCl exactly as shown in the boxes. Click on the ‘next’ button. 3. In the ‘Run experiment’ box, select the options for ‘show graph view’ and ‘show microscopic view’. Hit the start button to run the simulation. 4. Make note of the temperature after the simulation has completed. Record the initial (pre-set at 20 °C) and final temperatures for the trial. The specific heat and the density of the resulting solution can be approximated as those of pure water. Once you have noted all the numbers, click on the ‘Reset’ button on the bottom right corner to run the next experiment. Part 2 – The Molar Heat of Neutralization Molarity of HCl Molarity of NaOH Specific heat of solution Volume of NaOH solution (mL) Volume of HCl solution (mL) Initial solution temperature (°C) Final solution temperature (°C) Temperature change (?°C) Mass of solution (grams) Part 3 – The Molar Heat of Reaction 1. In the first column, click on the ‘solids’ tab. In the drop-down menu, select NaOH, and select an initial mass. Temperature will be pre-set at 20.0 °C. Record the mass and temperature of NaOH exactly as shown in the boxes. Click on the ‘next’ button in the panel on the left-hand side to proceed to the next column. 2. In the second column, click on the ‘solutions’ tab. Then, in the drop-down menu, select HCl (aq). Enter an initial volume and molarity of your choice. Record the volume and molarity of HCl exactly as shown in the boxes. Click on the ‘next’ button. 3. In the ‘Run experiment’ box, select the options for ‘show graph view’ and ‘show microscopic view’. Hit the start button to run the simulation. 4. The simulation will take about a minute to run. At the end of the experiment, a temperature will be displayed on top of the calorimeter. Make note of this temperature exactly as shown. Record the initial and final temperatures for the trial. Use the specific heat and density of pure water as the specific heat and density of the solution. Part 3 – The Molar Heat of Reaction Molarity of HCl Specific heat of solution Mass of NaOH pellets (grams) Volume of HCl solution (mL) Initial temperature of HCl (°C) Final solution temperature (°C) Temperature change (?°C) Mass of solution (grams) Once you have collected all the data, proceed to the worksheet. You can re-run any experiments that you would like to. Make sure to note all parameters used and the data collected. THE ENTHALPY OF REACTION AND HESS’S LAW Post -La b Works heet Write the balanced chemical reactions for the three processes you simulated in this experiment. Remember to include states of matter. Simulation Balanced chemical reactions 1 2 3 2. Examine the equations above. How are the processes from Part 1 and Part 2 related to Part 3? 3. Use the data collected in Part 1 to calculate the following: Show calculations and formula used Mol of NaOH Heat produced (q) in kJ Enthalpy Change (?H) in kJ/mol 4. Use the data collected in Part 2 to calculate the following: Show calculations and formulas used to calculate Mol of Limiting Reactant Heat produced (q) in kJ Enthalpy Change (?H) in kJ/mol 5. Use the data collected in Part 3 to calculate the following: Show calculations and formulas used. Mol of Limiting Reactant Heat produced (q) in kJ Enthalpy Change (?H) in kJ/mol 6.. Comparison of ?H values a. What is your expectation regarding the sum of the ?Hs for Parts 1 and 2 compared to the ?H for Part 3? Explain the reason for your expectation. b. Fill in the following: Sum of ?H from procedures 1 and 2 ?H from procedure 3 c. How does the sum of the ?Hs for Parts 1 and 2 compare to the ?H for Part 3? Does this meet your expectations? d. In a real-world experiment, how might you expect these values to differ? Why?