Heat of combustion

 

Introduction

 

The purpose of this experiment is to determine the enthalpy of combustion of an organic compound, and to determine its standard enthalpy of formation.  The apparatus to be used is a Parr 1341 Plain Jacket Oxygen Bomb Calorimeter.  See the instruction manual1, located in the lab, for a detailed description of the calorimeter.

Read this manual carefully before attempting an actual experiment. This experiment involves oxygen gas under high pressure.  Safety depends upon your following the instructions given in the Manual1.

 

The bomb calorimeter consists of a heavy-walled stainless steel reaction vessel containing a weighed sample of the compound to be burned and 25 atm of oxygen gas.  The high pressure of oxygen is necessary to insure complete combustion.  The combustion reaction is initiated by passing an electrical current through an iron fuse wire in contact with the sample.  The reaction vessel is immersed in an insulated water bath.   The heat released in the combustion reaction is absorbed by the water and the other parts of the calorimeter, thereby causing the temperature of the calorimeter to rise.  The calorimeter is adiabatic, that is, all of the heat liberated by the reaction remains in the calorimeter.  None is lost to the surroundings. The temperature rise of the calorimeter can be measured by a sensitive thermometer or other temperature measuring device.  We will use a calibrated thermistor connected to a Labworks interface to measure the temperature.  The reaction occurs under constant volume conditions, and hence the heat liberated by the reaction is equal to the internal energy change:

 

qRXN = ΔURXN

(1)

 

The actual process occurring in the calorimeter involves the conversion of reactants at some initial temperature, T1, to reactants at some final temperature, T2. Since the calorimeter is adiabatic, no heat is transferred between the calorimeter and the surroundings, and the internal energy change for the calorimeter as a whole is zero, that is DUCAL = 0.  However, heat transfers occur within the calorimeter. If the reaction is exothermic, heat is released by the reaction.  Some of this heat raises the temperature of the products and the rest is transferred to the water bath and the other parts of the calorimeter. The process can be described in terms of a thermodynamic cycle as follows:

 

                                               

 

 

 

 

Thus, since U is a state function:

 

ΔUCAL = 0 = ΔURXN + ΔUHEATING

and

 

ΔURXN = - DUHEATING =  

(2)

 

In the previous equation, C is the heat capacity of the calorimeter and products.  The mass of the products is negligible compared to the mass of the water in the bath and the other parts of the calorimeter, and their contribution to C is negligible.  C can be considered to be a constant property of the calorimeter.  Note that DURXN  is the internal energy change for the reaction at a constant temperature, T­1.

 

For the combustion of n moles of organic compound the molar internal energy of combustion at temperature T1, DUc,m is

 

 

 

The molar enthalpy of combustion, DHc,m is given by

 

 

where Dngas is the change  in the number of gaseous moles in the balanced chemical equation.  For example for the combustion of naphthalene:

 

  C10H8(s) + 12 O2(g) → 10 CO2(g) + 4 H2O(l)

 

Δngas = 10-12 = - 2 .

Although the molar enthalpy change depends on both temperature and pressure this dependence is not large, and we will neglect it here.  This allows us to equate the measured ΔHc,m to , the standard molar enthalpy of combustion at 298K.  For the combustion of naphthalene this leads to

DHc,m = ΔHoc(298) = 10 ΔfHo(CO2) + 4 ΔfHo(H2O) – ΔfHo(C10H8)

(3)

  where the DfHo terms are standard enthalpies of formation at 298 K.  From equation (3) it is evident that the standard enthalpy of formation of naphthalene can be calculated from its molar enthalpy of combustion and the standard enthalpies of formation of CO2(g) and H2O(l).

The heat capacity of the calorimeter, C, in equation (3), is determined by burning substance with a known DURXN.  High purity benzoic acid is commonly used for this purpose.  The reaction is

  C6H5CO2(s) + 15/2 O2(g) → 7 CO2(g) + 3 H2O(l)

The total heat liberated in the combustion process is due to three sources: the combustion of benzoic acid, the combustion of the fuse wire, and the formation of nitric acid resulting from the high temperature reaction of nitrogen from the air with oxygen and water.

The heat capacity of the calorimeter (J/oC) is calculated from the following equation:

 

(4)

Where

Qb = the heat of combustion of benzoic acid (26.43 kJ/g)

m = the mass of benzoic acid in g

e1 = correction for the heat of formation of nitric acid

e3 = correction for the heat of combustion of the iron fuse wire (5858 J/g)

DT = temperature rise accompanying the combustion of benzoic acid.

In this experiment, flushing the bomb with pure oxygen, before filling it, will eliminate the need for the correction term e1.

Once C is measured, the enthalpy of combustion (in J) is given by:

(6)

Determination of D T: The temperature of the calorimeter is measured as a function of time before, during and after ignition of the sample.  A graph of temperature versus time is called a thermogram.  An example is shown below.

Prepare the thermogram by plotting temperature versus time.  Perform a linear regression on the points in the pre-ignition period and those in the post-ignition period.  Plot the regression lines. Lines AB and CD, repectively. Perform a trendline on each line. Draw a line parallel to the y axis at time to such that Area E is approximately equal to Area F.  Using the value of to (here to = 7.1 min) calculate the initial and final temperatures, T1 and T2, respectively, using the trendline equations.  The temperature rise is D T = T2 – T1.

Figure X: A thermogram.

Procedure

      Calibration

1.    weigh about 0.9-1. g of benzoic acid into the pellet press. Press a pellet and place it in a pre-weighed metal sample cup. With the pellet in the cup, obtain the exact mass of the pellet to ± 0.1 mg.

2.    Place the bomb cover in its holder.

3.    Cut and weigh a 10 cm length of fuse wire.

4.    Attach the fuse wire to the electrodes as described in the Instruction manual.

5.    Place the sample cup and sample in the holder attached to the bomb cover. Make sure that the fuse wire is in good contact with the sample, and not touching any metal surface.

6.    Carefully lift the bomb cover from its holder, insert it into the lower part of the bomb, and tighten it by HAND only.

7.    Close the outlet valve. You are now ready to admit oxygen into the bomb.  DO NOT PROCEED WITH THIS STEP UNLESS THE INSTRUCTOR IS PRESENT. Attach the inlet tubing from the oxygen tank Make sure that the valve between the main valve of the oxygen tank and the bomb is closed.  Open the main valve on the oxygen tank.  Flush the bomb twice with 10 atm of oxygen.  Then fill it to a pressure of 25 atm.

8.    Place 2000 mL of room–temperature water, measured from a volumetric flask, into the stainless steel bucket. Place the bucket into the calorimeter.

9.    Using the special detachable handle, lower the bomb into the water bucket.  Be sure to attach the electrical leads to the fuse before the bomb is submerged.

10.Modify the labworks temperature collection program to collect the temperature every 30 s, and calibrate the thermistor. Carefully place the calorimeter lid in place.

11.Connect the stirrer motor to the stirrer with the drive belt and turn on the motor.  Wait at least 5 min.  Start collecting temperature-time data.

12. Collect data for 5 min, BEFORE igniting the sample. Ignite the sample by pressing the firing button for 5 s, AND NO LONGER. Within about 20 s of firing the temperature should start to rise.  Continue collecting data for 5 min after the temperature stops rising.

13. Remove the bomb from the calorimeter using the detachable handle.  Release the pressure and open the bomb.  Retrieve the unburned fuse wire, and weigh it.

14.  Repeat the calibration with two additional samples on benzoic acid.  That is, do the calibration in triplicate.

Measurement of enthalpy of combustion

Repeat steps 1 through 13 of the calibration procedure with the compound under investigation in place of benzoic acid.  Keep the sample size between 0.9 and 1 g. Perform two trials with this substance.

Calculations and Results

1.   Calculate the average heat capacity of the calorimeter, C, and the 95 % confidence limit.

2.      In equation (4), estimate the errors in m, e3 and DT, and estimate the probable propagated error in C. The errors can be propagated using the methods you learned in Quantitative Analysis.  An alternative method for the propagation of errors is described in this link.

3.      Calculate the average value of the molar internal energy of combustion, DUc,,m, and the 95 % confidence limit

4.      Calculate the probable propagated error in DUc,,m.

5.      Calculate the value of the molar enthalpy of combustion, and its probable error.

6.      Calculate the standard molar enthalpy of formation of the compound under investigation and the probable error.

Conclusion

7.      Report the measured value of the standard molar enthalpy of formation of the compound under investigation and the probable error. Compare your measured value with a literature value.

Reference

1.   Instructions for the 1341 Plain Jacket Oxygen Bomb Calorimeter, Parr Instrument Company, Moline IL, Manual No. 147.