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Inside Collection:

Collection by: Mary McHale. E-mail the author

# Thermochemistry

Module by: Mary McHale. E-mail the author

## Thermochemistry

We will begin lab in the amphitheatre of DBH

### Objective

• To explore heat transfer through calorimetry.
• To use calorimetry to determine the enthalpy of reaction of a strong acid and a strong base.
• To use Hess's Law of Heat Summation to determine the heat of hydration in calcium chloride.
• To explore a common use of heat of reaction in real life.

• Pre-Lab (10%)
• Lab Report Form (80%)
• TA Points (10%)

### Background Information

When two objects at different temperatures are brought into physical contact, thermal energy will spontaneously transfer from the warmer object to the colder object until both objects have achieved the same temperature. Assuming the two objects are thermally insulated from their surroundings, the heat lost by the warm object is identical to the heat gained by the cold object. This is a manifestation of the Law of Conservation of Energy.

The heat transfer, q, is a function of the mass of the object (m), the change in temperature undergone by the object ( ΔTΔT size 12{ΔT} {}) and the object's specific heat ( CsCs size 12{C rSub { size 8{s} } } {}). This statement can be expressed mathematically as

q = mCsΔTmCsΔT size 12{ ital "mC" rSub { size 8{s} } ΔT} {}

Temperature change is always defined as TfinalTfinal size 12{T rSub { size 8{ ital "final"} } } {} - TinitialTinitial size 12{T rSub { size 8{ ital "initial"} } } {}, which means that q for the hotter object ( qhotqhot size 12{q rSub { size 8{ ital "hot"} } } {}) is negative and q for the colder object ( qcoldqcold size 12{q rSub { size 8{ ital "cold"} } } {}) is positive. If energy is conserved, then

qhotqhot size 12{q rSub { size 8{ ital "hot"} } } {} + qcoldqcold size 12{q rSub { size 8{ ital "cold"} } } {} = 0

and

( mCsΔTmCsΔT size 12{ ital "mC" rSub { size 8{s} } ΔT} {} ) hothot size 12{ {} rSub { size 8{ ital "hot"} } } {} + ( mCsΔTmCsΔT size 12{ ital "mC" rSub { size 8{s} } ΔT} {} ) coldcold size 12{ {} rSub { size 8{ ital "cold"} } } {} = 0

Now consider dropping an ice cube into water just warm enough to melt the ice cube but not warm enough to further heat the water from the cube. The observation is that the ice cube melts and the warm water cools to 0°0° size 12{0°} {}C.  It is important to recognize that during the phase change, the temperature of the ice cube does not change. Therefore, it is not possible to use the preceding equation to determine the heat transferred.  Rather, the energy transferred to the ice cube from the warm water affects the phase change. The energy equation is now adjusted to incorporate the enthalpy required to melt the ice cube, ΔHfΔHf size 12{ΔH rSub { size 8{f} } } {} (where f stands for fusion):

( mCsΔTmCsΔT size 12{ ital "mC" rSub { size 8{s} } ΔT} {} ) warmwarm size 12{ {} rSub { size 8{ ital "warm"} } } {} + ΔHfΔHf size 12{ΔH rSub { size 8{f} } } {} = 0

It is also possible to have thermal energy when chemical reactions occur. The amount and direction of heat flow is dependant on the chemicals reacting. Using a calorimeter, it is possible to experimentally determine the heat of reaction.

#### Calorimetry

In the technique known as constant-pressure calorimetry, enthalpies of phase changes or chemical reactions are determined indirectly by measuring temperature (at constant pressure) changes in a medium, most often water, surrounding the materials undergoing the change. That is, by measuring ΔTΔT size 12{ΔT} {} of the water one can use the preceding equation to calculate ΔHΔH size 12{ΔH} {} for the process of interest.  Of course, this means one must know the mass of the water used and water's specific heat: CwaterCwater size 12{C rSub { size 8{ ital "water"} } } {} = 4.18 J/(gK).

Today in Part I, you will add a strong base to a strong acid, measure the temperature change in the water as the two react, and use that information to calculate the heat of reaction per gram of NaOH. Then convert your experimental value into an enthalpy in kJ/mol (of NaOH).

The enthalpy of a reaction might be difficult to obtain directly by experiment but can be determined by measuring (or looking up in tables) the enthalpy changes of reactions which contain the reactants and products in a process governed by Hess's Law of Heat Summation. In Part II of this experiment you are asked to find the enthalpy of hydration of CaCl2(s)CaCl2(s) size 12{ ital "CaCl" rSub { size 8{2} } $$s$$ } {} the "target reaction" in this manner.

Target reaction:  CaCl2+6H2O=CaCl26H2OCaCl2+6H2O=CaCl26H2O size 12{ ital "CaCl" rSub { size 8{2} } +6H rSub { size 8{2} } O= ital "CaCl" rSub { size 8{2} } cdot 6H rSub { size 8{2} } O} {}

It cannot be measured directly because of the slow kinetics of the reaction in the solid state. However, the heats of dissolution of CaCl2CaCl2 size 12{ ital "CaCl" rSub { size 8{2} } } {} and CaCl26H2OCaCl26H2O size 12{ ital "CaCl" rSub { size 8{2} } cdot 6H rSub { size 8{2} } O} {} can be determined and, using Hess's Law, the enthalpy of hydration can be calculated.

Theoretical value of the enthalpy of hydration of CaCl2CaCl2 size 12{ ital "CaCl" rSub { size 8{2} } } {} anhydrous is -81.33kJ/mol. Theoretical value of the enthalpy of hydration of CaCl26H2OCaCl26H2O size 12{ ital "CaCl" rSub { size 8{2} } cdot 6H rSub { size 8{2} } O} {} is 15.79 kJ/mol.

## Experimental Procedure

### Materials Required

• two Styrofoam cups (serve as the calorimeter)
• styrofoam cover (top to the calorimeter)
• conventional or digital thermometer
• stirring motor and stir bar
• two large Ziploc bag
• two plastic spoons

#### Setup of the MicroLab Thermistor

• Open the MicroLab Program by clicking on the Shortcut to MicroLab.exe tab on the desktop.
• On the “Choose an Experiment Type” Tab, enter a name for the experiment, and then double click on the MicroLab Experiment icon
• Click “Add Sensor”, Choose sensor = Temperature (thermistor)
• Choose an input, click on the red box that corresponds to the port that your thermistor is connected to.
• Label = Thermistor, sensor units = °C°C size 12{°C} {}, click next
• Click “Perform New Calibration”
• Click “Add Calibration Point” place the thermistor and thermometer in an ice water bath. Wait until the temperature is constant and then read the temperature on the thermometer (to the nearest 0.1 °C°C size 12{°C} {}) and enter it into the “Actual Value” box in MicroLab and hit “ok”.
• Again, Click “Add Calibration Point” place the thermistor and thermometer in warm water bath. Wait until the temperature is constant and then read the temperature on the thermometer (to the nearest 0.1 °C°C size 12{°C} {}) and enter it into the “Actual Value” box in MicroLab and hit “ok”.
• Under Curve Fit Choices , click on “First order (linear)” and then “Accept and Save this Calibration”, when prompted to enter units, set as deg C. Save as your name-experiment-date.
• Click “Add Sensor”, Choose sensor = Time
• Choose an input, click on the red box that corresponds any of the Timers.
• Label = Time 1, click next, click Finish.
• Left click on thermistor and drag to the Y-axis over “data source two”, also click and drag to column B on the spreadsheet and also click and drag to the digital display window.
• Left click on Time and drag to the X-axis over “data source one”, also click and drag to column A on the spreadsheet and also click and drag to the digital display window.
• When you are ready to obtain data, click start. When you are finished collecting data, click stop. To run another trial, click repeat experiment.

### Part I. Reaction of Strong Acid with Strong Base

1. Weigh the calorimeter (styrofoam cups with lid) WITH a stir bar. Accurately measure 50 mL of 1 M HCl into your calorimeter. Turn on the stirring motor to the medium speed. Fit the thermistor though the cardboard lid to a length such that its tip goes deep into the solution but misses the stir bar. Cover the calorimeter.
2. Start the Acquisition program. You should see successive, constant temperature readings of near room temperature. Accurately measure 50 mL of 1 M NaOH and quickly add it to the acid solution in the Styrofoam cups. Cover the calorimeter.
3. Continue monitoring the temperature change until thermal equilibrium is reached (the temperature stops changing or starts decreasing).
4. Stop the Acquisition program. SAVE YOUR FILE. Remove the cardboard lid, reweigh the calorimeter and record the information on your lab report.
5. Repeat the same procedure (Steps 1-4) two more times. Save all your files. Give them different, distinguishable names.

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°C size 12{°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

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

1. 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.
2. 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.

1. Continue monitoring the temperature change until thermal equilibrium is reached (the temperature stops changing or starts decreasing).
2. Stop collecting and SAVE YOUR DATA and/or print it.
3. 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.
4. Repeat the process but replace anhydrous with .
5. 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.

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