0.1 Thermochemistry  (Page 3/3)

Print off your graphs and use the data to determine the initial and final temperatures. Record all the temperatures to the nearest 0.1 $°C$ . If a slight downward trend appears on the final temperature plateau, use the maximum value achieved. Calculate the average enthalpy of the reaction and standard deviation for the three trials. Then combine your data with the data obtained from the rest of the group and calculate average enthalpy of the reaction and standard deviation for all trials together. Compare your average and standard deviation with that of the larger set and comment on the results obtained using a larger data set.

Part ii. enthalpy of reaction

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The same basic procedure you used in the dry ice sublimation lab will be used here today. 

• Tare your calorimeter and pour approximately 75 mL of cold water in it. Add stir bar. Record the mass to 3 decimal places. Place the cups on the magnetic stirrer and turn on the stirring motor to a medium rate. Make sure you can fit the thermistor though the styrofoam cover to a length such that its tip goes deep into the water but misses the stir bar but do not insert it. Cover the calorimeter.
• Start recording data on the MicroLab interface. Let it run for about 10 seconds before putting thermistor through the Styrofoam cover. You should see successive, constant temperature readings of near room temperature. Approximately 5 grams of anhydrous has been weighed for you, record the weight of this plus the weight of the weighing bottle. Then weigh the dry empty weighing bottle. Quickly add to the water and reposition the thermistor and cover assembly.

Note: Since anhydrous absorbs moisture rapidly from the air, close the lids of the bottles securely immediately after using. Dry the spatula each time before weighing the powder and clean the balance of any solid.

• Continue monitoring the temperature change until thermal equilibrium is reached (the temperature stops changing or starts decreasing).
• Stop collecting and SAVE YOUR DATA and/or print it.
• Repeat steps 1-4 two more times. Don’t forget to save your data each time you do a run because it is lost as soon as the next run begins.
• Repeat the process but replace anhydrous with .
• Use the data to determine the initial and final temperatures. Record all temperatures to the nearest 0.1 . If a slight downward trend appears on the final temperature plateau, use the maximum value achieved. Calculate the average enthalpy of dissolution and standard deviation. Use Hess’s Law to calculate the heat of hydration of and % error. Then combine your data with the data obtained from the rest of the class and calculate the average enthalpy of dissolution and hydration.

Calculations

Calculations are similar to those done for the acid-base neutralization reaction. The calculation of is the same as the calculation of , that is:

=

qwater can be calculated using water’s specific heat, mass of water and temperature change of water solution:

In order to account for the mass of anhydrous or hexahydrate, divide by the mass of anhydrous or to get in J/g. Then convert to J/mol.

Part iii. chemistry of life

Hot packs and cold packs are a real life example of thermochemistry. Anhydrous magnesium sulfate and ammonium chloride can be used to make“hot/cold”packs similar to those used for sports injuries and in hospitals. Your TA will make a pack from each of the two compounds and pass them around and answer some fundamental thermochemical questions about the reactions involved.

Ta procedure

• Fill two Ziploc bags half way full with water.
• Put two spoons full of anhydrous magnesium sulfate into one bag and two spoons full of ammonium chloride into the other bag, zip them closed, and shake.

Pass the packs around to students and observe the temperature change.