Why is polystyrene used to construct a calorimeter?

Polystyrene is an excellent insulator that minimizes heat exchange between the calorimeter's interior and external environment. This property ensures that the heat measured during reactions comes primarily from the chemical reaction itself rather than from heat loss to surroundings, making temperature measurements more accurate and reliable for calculating enthalpy changes.

What is the purpose of calibrating a calorimeter before measuring reaction heat?

Calorimeter calibration determines how much heat the calorimeter itself absorbs during temperature changes. By adding hot water to cold water and measuring the temperature change, you calculate the calorimeter's heat capacity. This value is then used to account for heat absorbed by the calorimeter during chemical reactions, ensuring accurate enthalpy measurements.

How does polishing magnesium ribbon affect the reaction with hydrochloric acid?

Polishing removes the outer layer of magnesium oxide that forms when magnesium metal is exposed to air. This reveals fresh, reactive magnesium metal underneath, ensuring the reaction proceeds efficiently and completely. Without polishing, the oxide layer would slow or inhibit contact between the magnesium and hydrochloric acid.

Why is stirring important during calorimetry experiments?

Stirring ensures uniform mixing of reactants and products throughout the solution, allowing the temperature probe to accurately measure the true temperature of the entire mixture. Vigorous stirring also helps dissolve reactants completely and distributes heat evenly, preventing localized hot or cold spots that would compromise temperature readings and enthalpy calculations.

What assumptions allow enthalpy change to equal internal energy change in these experiments?

The experiments assume constant pressure, negligible volume change as the solution heats, and no energy transfer between the solution and surroundings. Under these conditions, the enthalpy change equals the heat released or absorbed, which approximately equals the change in internal energy. These assumptions simplify calculations and make the measured temperature changes directly proportional to reaction enthalpy.

How do you calculate the enthalpy of formation of magnesium oxide from two separate reactions?

You measure the enthalpy of the magnesium-hydrochloric acid reaction and the magnesium oxide-hydrochloric acid reaction separately. Then you reverse the second reaction and combine it with the first reaction and the known enthalpy of formation of water using Hess's Law. This combination yields the enthalpy of formation of magnesium oxide, which you compare to literature values to determine percent error.

Why must magnesium oxide and hydrochloric acid reaction data be collected after the magnesium metal reaction?

The second reaction is performed after the first to allow time for data processing and to ensure the calorimeter is thoroughly cleaned and dried between experiments. Additionally, the experimental design uses both reactions together with Hess's Law to calculate the enthalpy of formation of magnesium oxide, requiring both datasets to be complete before final calculations can be performed.

Enthalpy of Reaction

Concepts Instructor Prep Student Protocol

0 views03:37 min

Thermodynamics is the study of heat and its relation to energy and work. One of the most important thermodynamic properties is enthalpy, which indicates the heat energy transferred by a process at constant pressure. Enthalpy is used to describe chemical reactions, where the enthalpy change, ΔH, tells us how much heat is absorbed or released during a chemical reaction.

Enthalpy is a state function, meaning that the change is independent of the path and only takes into account the initial and final state. Thus, the ΔH of a reaction is calculated by subtracting the sum of the enthalpies of the reactants from the sum of the enthalpies of the products.

If ΔH is positive, the reaction is endothermic and absorbs heat from the surroundings. If ΔH is negative, the reaction is exothermic and releases heat to the surroundings.

We can measure the enthalpy change of a reaction using an isolated system like a calorimeter. A simplified calorimeter, like the one we have here, consists of an insulated sample cell with a stirrer and a thermometer. As the reaction proceeds inside of the insulated cell, the temperature changes and we are able to measure ΔT.

We can assume that no energy is transferred into or out of the sample cell because it is insulated. So, we can then assume that the temperature change is due to the heat change in the reaction.

So, how do we use ΔT to calculate ΔH? At constant pressure, ΔH is equal to the heat flow, Q. Q is equal to the mass of the substance times its specific heat capacity and the change in temperature. The specific heat capacity, c_s, is a measure of how much heat is required to increase the temperature of 1 g of a material by 1 °C.

A material with a higher specific heat capacity requires more energy to raise its temperature and vice versa. Specific heat capacity values for common substances can be found in the literature. Thus, we can use the specific heat capacity, along with the mass of the substance and the temperature change, to calculate ΔH for the reaction.

In your experiment, we will measure the change in enthalpy for the formation of magnesium oxide from magnesium and oxygen. Using our simple calorimeter requires us to measure the heat of reaction indirectly with a series of separate reaction steps. So, we will run each reaction separately and measure ΔT and, by extension, ΔH for each reaction. Then, according to Hess's Law, the enthalpy of the overall reaction is equal to the sum of each of the enthalpies of reaction for each reaction step.

In this lab, you will construct a simple calorimeter using polystyrene cups and then experimentally determine the enthalpy of formation of magnesium oxide.

Print Procedure Steps

$$\rightleftharpoonup{xx}$$ $$\longleftharp{xx}$$, $$\longrightharp{xx}$$,

Source: Smaa Koraym at Johns Hopkins University, MD, USA

Collecting Data for Calculating the Heat Capacity of a Calorimeter

This experiment will use a constant-pressure calorimeter made of two stacked polystyrene cups, a cardboard lid, and a probe measuring the temperature of a solution in real-time. Polystyrene is a good insulator, so you can assume that no heat is exchanged between the inside and outside of the cups. The calorimeter itself will absorb some heat, which you'll account for by calculating its heat capacity.

In the first part of this experiment, you'll add hot water to cold water in the calorimeter and measure the temperature increase. After the experiment, you'll calculate how much heat the hot water lost to the calorimeter, rather than to the cold water.

Put on a lab coat, splash-resistant safety glasses or goggles, and disposable gloves. Make sure that your fume hood is clean and ready to use and that your glassware is clean and undamaged.
To begin the lab, stack two polystyrene cups to form the insulated calorimeter vessel. Place a magnetic stir bar in the calorimeter.

Measure the combined mass of the empty cups and the stir bar on a zeroed top-loading balance. Record this in your lab notebook as the empty calorimeter mass and return to your fume hood.

Table 1: Calorimeter Calibration

Mass_{empty calorimeter}(g)
Mass_{calorimeter + cold water} (g)
Mass_{cold water} (g)
T_{cold initial} (ºC)
T_{hot initial}(ºC)
Mass_{calorimeter + cold and hot water mix}(g)
Mass_{hot water} (g)
T_final(ºC)
ΔT_hot (K)
ΔT_cold (K)
Specific heat capacity of water (J/g·K)	4.184
Heat capacity of calorimeter (J/K)

Click Here to download Table 1

Measure 50 mL of cold tap water and pour it into the calorimeter. Measure the combined mass, record it in your lab notebook, and bring the calorimeter back to your workstation.
Calculate the mass of the cold water, or m_cold, by subtracting the mass of the empty calorimeter from the mass of the calorimeter with cold water in it.
Measure 50 mL of deionized water and pour it into a 150-mL beaker. Place the beaker on a hot plate under a thermometer clamp and fix a digital thermometer in the water.
Set the calorimeter in the center of a small stir plate under another thermometer clamp and cover it with a cardboard lid that has a hole in the center.
Turn on your data-recording device and configure it to record the temperature over time. Set the data-recording rate to two readings per second and the recording duration to 600 s.
Insert the temperature probe through the hole in the calorimeter lid and clamp it in place. The probe should always be fixed high enough that you can lift the lid.
Start the stir motor, and increase the rotational speed until the stir bar is stirring steadily in the calorimeter.
Once the temperature shown on the data acquisition device stabilizes to within plus or minus 0.1 °C, record it in your lab notebook as T_cold-initial.
Then, turn on the hotplate and monitor the water temperature with the digital thermometer. While you wait, bring some paper towels to your fume hood.
Once the water reaches 60 °C, turn off the hot plate. Remove the temperature probe from the cold water, wipe it off, and hold it or clamp it in the hot water. Ensure that the probe does not touch the bottom of the beaker as the temperature reading stabilizes.
Record the temperature in your lab notebook as T_hot-initial. Then, remove the probe from the hot water, wipe it off, and return it to the calorimeter.
Once the temperature reading has stabilized, start data collection and wait for the device to record a few seconds of data.
Then, carefully pour the hot water into the calorimeter. Monitor the temperature reading as the hot and cold water mix.
Once the temperature reaches a maximum, meaning that the temperature is no longer increasing, continue collecting data for 5 - 6 min to ensure that you record an equilibrium temperature.
Then, stop data collection and name the data ‘calibration’. Remove the probe and lid from the calorimeter and set them aside.
Measure the mass of the water-filled calorimeter on the top-loading balance.
Record the mass in your lab notebook and bring the calorimeter back to the fume hood.
Calculate the mass of the hot water, or m_hot, by subtracting the mass of the calorimeter containing cold water from the mass of the calorimeter containing mixed hot and cold water.
Remove the magnetic stir bar from the calorimeter and pour the water down the sink.
Thoroughly dry the calorimeter and the stir bar and return the stir bar to the calorimeter before proceeding.

Mg_(s)+ 2 HCl_(aq) → MgCl_2(aq)+ H₂_(g)

This is the first reaction that you will perform in the calorimeter. Here, magnesium metal is oxidized to magnesium 2+ and H+ is reduced to hydrogen gas. You will continue working in a fume hood because of the hazards associated with the reactants and products of this reaction. Magnesium metal ribbon is water-reactive and flammable. Hydrochloric acid is corrosive and toxic, and hydrogen gas is also flammable.

Create a new data collection run with the same settings as before.
Label a 600-mL beaker for aqueous magnesium waste, which must be disposed of separately at the end of the lab.
Obtain the magnesium ribbon, which is provided in 0.45-to 0.55-g pre-cut strips. Place one piece of ribbon in a weighing boat and bring it back to your hood.
Then, use a paper towel and emery paper to polish the magnesium ribbon, revealing shiny, silvery material. Polishing removes the outer layer of magnesium oxide that forms when magnesium metal is exposed to air.
Cut the polished magnesium ribbon into 0.5 to 1-cm long pieces in the weighing boat.

Use an analytical balance and a second tared weighing boat to measure the precise mass of your magnesium. Record the mass in your lab notebook. Then, discard the extra weighing boat and bring the magnesium pieces back to the fume hood.

Table 2: First reaction Mg + 2 HCl

Mass_Mg (g)
Volume_HCL _added (mL)
T_initial (°C)
T_final (°C)
ΔT_{Mg + HCl} (K)
Density_HCl (g/mL)	1.039
Mass_HCl (g)
Mass_solution(g)
Specific heat capacity of 2 M HCl (J/g·K)	3.98
ΔU_{Mg + 2H}+≈ ΔH_{Mg + 2H}+ (kJ)
Limiting reagent (mol)
Theoretical yield (mol)
ΔH_{Mg + 2H}+(kJ/mol)

Click Here to download Table 2

Now, bring a 150-mL beaker and a watch glass to the fume hood for dispensing HCl, and carefully pour about 110 mL of 2.0 M HCl into the beaker. Cover the beaker and transport it back to your hood.
Use a 50-mL graduated cylinder to add 100 mL of 2 M HCl to the dry calorimeter. Record the total volume of HCl added and neutralize any drops of spilled acid.
Then, put the calorimeter on the stir plate, cover the calorimeter, and clamp the temperature probe in place.
Set the stir motor to a medium-high speed and start recording temperature data. Wait for 30 s of data to be collected so that you obtain an accurate initial temperature.
Then, lift the cardboard lid and add the magnesium pieces to the calorimeter all at once.
Adjust the stir speed so that the pieces are stirring briskly. Note: It's important to stir the mixture well to dissolve as much magnesium in the acid as possible.
Now, close the calorimeter and monitor the data readout as the temperature increases. Once the temperature reaches a maximum, allow data collection to continue for another five minutes. Then, stop data collection and turn off the stir motor.
To avoid cross-contamination in your next experiment, thoroughly rinse the temperature probe with deionized water and collect the rinse in the waste beaker.
Then, remove the calorimeter from the stir plate and retrieve the stir bar with forceps. Clean the stir bar and forceps in the same way.
Pour the contents of the calorimeter into the waste and thoroughly wash the calorimeter with deionized water. Dry the calorimeter with paper towels.
Switch the stacking order of the polystyrene cups to provide a clean inner surface for the next reaction. Dry the forceps, stir bar, and temperature probe and return the stir bar to the calorimeter.

MgO_(s)+ 2 HCl_(aq) → MgCl_2(aq)+ H₂O_(l)

Here is the second reaction that you will perform in the calorimeter. This is an acid-base reaction in which magnesium oxide neutralizes HCl. As before, magnesium chloride will remain in solution as solvated ions. Magnesium oxide is a respiratory irritant, so use caution when measuring it out.

Create a new data collection run with the same settings as before to continue with the lab.

Measure out approximately 1 g of magnesium oxide on an analytical balance and record the precise mass in your lab notebook.

Table 3: Second reaction MgO + HCl

Mass_MgO (g)
Volume_{HCL added} (mL)
T_initial(ºC)
T_final (ºC)
ΔT_{MgO + HCl}(K)
Density_HCl (g/mL)	1.039
Mass_HCl (g)
Mass_Solution (g)
Specific heat capacity of 2 M HCl (J/g·K)	3.98
ΔU_{MgO + 2H}+≈ ΔH_{MgO + 2H}+(kJ)
Limiting reagent (mol)
Theoretical yield (mol)
ΔH_{MgO + 2H}+(kJ/mol)

Click Here to download Table 3

Use your 150-mL beaker to transport about 110 mL of 2 M HCl to your fume hood. Pour 100 mL of HCl into the calorimeter and record the HCl volume in your lab notebook.
Set up the calorimeter on the stir plate in the same way as the last experiment. Begin stirring at high speed and start collecting temperature data. Wait 30 s to establish an accurate initial temperature reading.
Then, lift the calorimeter lid and pour the magnesium oxide into the calorimeter all at once. Ensure that the mixture is stirring vigorously before replacing the lid. Monitor the temperature reading as it increases.
Once the temperature reaches a maximum, collect 5 more min of data. Then, stop data collection and turn off the stir motor.
Save your data and transfer the file to a computer or flash drive, so you can process it later.
Now, clean and dry the probe, stir bar, and forceps in the same way as before. Pour the reaction mixture into the waste beaker and rinse the inside of the calorimeter before you dispose of the cups.
Neutralize the acidic magnesium waste with baking soda, confirm that the pH is acceptable with pH paper, and pour it into a container designated for aqueous magnesium waste.
Thoroughly rinse the beaker with deionized water and pour the rinse into the waste. Close the waste container when you are finished.
Then, put away your lab equipment and wash your glassware according to your lab standard procedures.
Neutralize any non-magnesium-containing aqueous acidic waste with baking soda before flushing it down the drain.
Discard the weighing boats, emery paper, and cups appropriately for your lab, clean up spilled materials in your fume hood, and throw out used paper towels and other trash.

Results
- First, plot your three sets of data with temperature on the y-axis and time on the x-axis.
- Find the changes in temperature that you need for your calculations, starting with your calibration data. You measured the initial hot and cold temperatures directly, so you only need to find T_final, which is the peak stable temperature after the water mixed. Fill in the initial and final values and solve for ΔT_hot and ΔT_cold.
- Next, look at the data from your reactions. As with the calibration data, Tfinal for each reaction is the peak temperature. T_initial is the stable temperature before each reaction began. Fill in these values and solve to get ΔT for each reaction. Since ΔT is a temperature difference and degrees Celsius and Kelvin have the same interval, you can just change the units to Kelvin.
- Now, calculate the heat capacity of your calorimeter and the changes in internal energy of the reactions. For these calculations, assume that the pressure was constant, the volume of the solution barely changed as it heated, and that there was no transfer of energy between the solution and its surroundings. Under these assumptions, the enthalpy change is equal to the heat, which in turn is approximately equal to the change in internal energy.
- First, calculate the heat capacity of your calorimeter. Start by filling in the masses of the hot and cold portions of water from the first part of the lab. Fill in the ΔT_hot and ΔT_cold values that you calculated from your calibration data. Use the specific heat capacity of water at ambient pressure and solve for the heat capacity of your calorimeter.
- Next, calculate the internal energy of the first reaction. Determine the mass of your solution by adding together the masses of your reactants. ΔT comes from your graph. You know the heat capacity of your calorimeter, and the specific heat capacity of 2 M HCl is about 3.98 J/K⋅g. Fill in these values, and solve to find the internal energy, and thus the approximate enthalpy of this reaction. Use the same process to estimate the enthalpy of the second reaction.
- For each reaction, identify the limiting reagent and calculate the theoretical yield in moles. Based on this, scale the enthalpies for a yield of 1 mole.
- Now, calculate the enthalpy of formation of magnesium oxide using these values and the known enthalpy of formation of H₂O. Note that your second reaction must be reversed for these three reactions to add up to the formation of magnesium oxide, so you'll use the negative entropy of the forward reaction.
  Table 4: Calculating the enthalpy of formation of MgO
  
  ΔH_{Mg + 2H}+(kJ/mol)
  -ΔH_{MgO + 2H}+(kJ/mol)
  ΔH_fH₂O (kJ/mol)
  ΔH_MgO (calculated value; kJ/mol)
  ΔH_MgO (literature value; kJ/mol)
  Percent error
  Click Here to download Table 4
- Finally, calculate the percent error between your calculated value and the known enthalpy formation of magnesium oxide.