Thermochemistry and Hess's Law

Purpose
To measure the enthalpy change of two different reactions in the laboratory.
To use Hess's Law to estimate the enthalpy change for the reaction: 2 Mg (s) + O₂ (g) 2 MgO (s)

Introduction
In this lab, you will carry out the following two reactions to determine the enthalpy change for each:

(a)      Mg (s) + 2 HCl (aq) MgCl₂ (aq) + H₂ (g); ΔH = measured in lab

(b)      MgO (s) + 2 HCl (aq) MgCl₂ (aq) + H₂O (l); ΔH = measured in lab

You will then use the two equations and their enthalpy changes, along with the thermochemical equation:

(c)      2 H₂ (g) + O₂ (g) 2 H₂O (l); ΔH = -571.6 kJ

and Hess's Law of Heat Summation in order to predict the enthalpy change of the reaction:

(d)      2 Mg (s) + O₂ (g) 2 MgO (s); ΔH = ?

In order to perform the calculations at the end of this lab, there are a few things you must know:

1. The heat capacity (C) of the coffee cup calorimeter is given as 10 J/°C.
        

2. The specific heat (s) of the solution will be estimated to be the same as that of water, or 4.18 J/(g·°C).
        

3. The systems you are studying are reactions (a) and (b) above. The surroundings include the calorimeter and the solution.
        

4. Use the combined mass of all reagents (solids and solutions) used in each trial for the mass of the solution (m).
        

5. The change in temperature (Δt) is determined by t_max - t_min.
        

6. q = C x Δt and q = s x m x Δt.
        

7. Heat gained by the system is lost from the surroundings. Heat lost by the system is released to the surroundings. This is otherwise known as the Law of Conservation of Energy. (q_system = -q_surroundings)
        

8. The enthalpy change for each reaction is equal to the heat of reaction at a constant pressure. (ΔH = q_system).
        

9. The enthalpy change determined in the previous step must first be converted to kJ per mole (kJ/mol) of the limiting reactant, then converted to represent the number of moles of the limiting reactant in the balanced equation. Seek assistance from the instructor, if necessary.

Materials

Procedure
You will work in the same groups (of 3-4 students) as were assigned in the magnesium/hydrochloric acid lab. Students who served as a manager or computer operator will work as lab techs for this activity. In the event you had only 3 students the first time, or if a group member has dropped since then, you may make your own assignments, but you should take on a different role than last time. However, each individual should have had a chance to do some of the dirty work, as well as some of the greater responsibility of getting the assignment submitted to the professor in a timely manner. To refresh your memory on the role played by each person, you may check out the group work page.

1. Plug a stainless steel temperature probe into the the LabPro®  interface and launch the Logger Pro application. Adjust the experiment length to 200 s and set the sampling rate to 5 seconds per sample. Adjust the number of decimal places for the temperature in your data table to the nearest ±0.1°C. Also, adjust the axes on your graph to accommodate a temperature range of 10-50°C.
        

2. Place your calorimeter onto an electronic balance (use a cheap, less precise balance for this measurement, as you do not want to risk spilling acid on the analytical balance) and tare the balance. Remove the calorimeter from the balance and carefully add 25.0 mL of 1.0 M HCl. Place the cup back onto the balance and record the mass of HCl added. Also, record the volume of HCl used.
        

3. Measure and record the mass of a magnesium strip (~0.15 g), using an analytical balance.
        

4. Slide the temperature probe into the small notch cutout on the cardboard cover and place the probe into the HCl. Stir the HCl with the probe to maintain a uniform temperature throughout the solution. Wait until the temperature stabilizes.
        

5. Roll the magnesium ribbon into a loose ball. Click the green "Collect" button to begin data collection. After a couple of data points have been collected, slide the cover aside and drop the ball of magnesium into the calorimeter. Slide the cover back into place. Continue stirring until the data collection ends.
        

6. Select Store Latest Run from the Experiment menu and save your data to disk.
        

7. Repeat Steps 2-6 twice to obtain a total of three trials with the magnesium reaction with hydrochloric acid.
        

8. Perform Steps 2-6 three times using ~0.25 g of magnesium oxide in place of the magnesium strips.
        

9. If your data looks good, copy and paste the data from each of your trials into an Excel spreadsheet. Use the appropriate function or formula to determine the minimum and maximum temperatures reached in each trial.

Assignment (use the answers to questions 1-3 in your Results and to questions 4-7 in your Discussion)

1. Determine which substance is the limiting reactant in the reaction of magnesium with hydrochloric acid, then in the reaction of magnesium oxide with hydrochloric acid. How many moles of the limiting reactant were used in each of the reactions? Provide at least one sample calculation for each of the two types of reactions.
        

2. Calculate the enthalpy change (ΔH) per mole of the limiting reactant, in kJ/mol, for each of the two reactions. Provide at least one sample calculation for each of the two types of reactions. (If you are submitting the report electronically, you may place formulas in spreadsheet cells instead.) See the introduction for helpful information and equations.
        

3. Then determine ΔH, in kJ, for the balanced equations in (a) and (b). (Look at the coefficients of the limiting reactant in each of the equations. Don't make this conversion difficult.) Calculate the average ΔH value for each of the two reactions and  submit your average ΔH values here before leaving lab (or within 24 hours, with instructor approval).
        

4. Use Hess's Law of Heat Summation, the class average enthalpy changes for reactions (a) and (b), and the enthalpy change for (c) given above to determine the enthalpy change for reaction (d). Show your work.
        

5. Calculate the percent error in the ΔH you calculated for reaction (d), assuming the reaction was carried out under standard thermodynamic conditions. (Use your calculated ΔH for reaction (d) and the known ΔH for the same reaction using the values found in Appendix IIB of the Tro textbook).
       

6. What would happen to the value of ΔH you calculated if all of your temperature readings were too high by 1°C?
        

7. What were some possible sources of error in this experiment? Explain.

Laboratory Report
You will be turning in a group laboratory report.  The report should include the title information, an introduction to thermochemistry in general (different from the introduction in this lab manual), experimental details (properly referenced if you choose that route), results, discussion, and references.  You should use the above questions to guide your results and discussion, but there should be more to these sections than just answering the questions and they show flow logically as you discuss the lab and the results.

Follow your instructor's directions for submitting this lab report. If you are submitting electronically, please submit a single file that includes all of the required components. Also, use the following convention for naming your files: Lastname1 Lastname2 Etc Thermo and if emailing, use a subject line of Chem 1061: Thermochemistry Lab.  As alluded to earlier, you do need to show calculations.  If you email the report, the embedded data tables must contain all formulas in calculated cells.  If you submit a paper version, you will need to show these calculations for one of the the trials.  In both cases, any calculations not carried out in the tables must be shown.  Laboratory report guidelines are found at http://webs.anokaramsey.edu/chemistry/Chem1061.

Written by Lance S. Lund 2000 (updated June 14, 2011).