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Heats of Combustion
Teaching Assistant: TA Name
Performed: 10/31/96, 11/7/96
Lab Partner: Their Name
A Parr bomb style calorimeter with an air filled adiabatic jacked was used to determine the identity of an unknown substance. The calorimeter constant was determined by igniting samples of benzoic acid with a piece of burning iron wire via an electronic ignition switch. Taking into account the heat introduced by the burning wire and the electrical charge, the calorimeter constant was determined and used to determine the heat of combustion of our unknown.
Introduction & Theory
Chemical bonds are one way that a system stores energy. As the bonds in a system are created or destroyed, energy is taken in or given off, usually in the form of heat. Different bonds store differing amounts of energy, so chemical reactions often absorb or release heat energy, even if the number of bonds created is equal to the number of bonds destroyed.
The energy of a bond depends not only the two atoms involved, but also other bonds and atoms in the molecule. As a result, the C-H bonds in organic molecules may vary widely depending on if the carbon is also attached to three flourine atoms, or perhaps three other hydrogens. Naturally, carbon to hydrogen bonds are by no means the only bonds affected.
Although bonds within a molecule may differ, the sum of bond energies within a given molecule is generally a constant (neglecting spatial considerations, which would average out in samples of practical size). As a result, a given substance possesses a certain amount of energy in each molecule, and accordingly in each mole.
As substances combust, some of the bonds of their molecules are broken and other bonds are formed, resulting in a release of heat energy. The amount of energy released in this fashion per mole of substance is referred to as the heat of combustion for a substance.
The procedure on pages 154-156 of Experiments in Physical Chemistry(1) was followed, with two runs of benzoic acid on 10/31 and one run of benzoic acid on 11/7 and three runs of unknown A.
The only exceptions to the procedure listed was in the way the fuse wire was attached to the pellet. Rather than melting it on as is suggested, a safer method of simply running the wire across the pellet was employed. Also, the pressure in the oxygen cylinder was low, preventing pressurization up to a full 25 atm on the first day. There was no evidence of soot however, so the amount of oxygen present was apparently still sufficient, so the results should not have been affected.
On 10/31 the ambient conditions were not recorded, and on 11/7 the ambient temperature was 23.1 C and the temperature corrected barometric pressure was 743.6 mm Hg. The lack of readings on the first day of the experiment is of no consequence as the calorimeter is designed to be adiabatic and hence, independent of external temperatures. Barometric pressure, so long as it is within reasonable ranges, has no effect on measurements.
The following tables of data were obtained from the experiment:
|Compound||Benzoic Acid||Unknown A|
|Initial wire mass (g)||0.0161||0.0159||0.0131||0.0156||0.0161||0.0155|
|Pellet mass (g)||0.8477||0.7848||0.8731||0.5382||0.5743||0.5971|
|Final wire mass (g)||0.0058||0.0052||0.0051||0.0047||0.0070||0.0051|
|Temperature (C) 0:00||21.50||21.11||22.88||23.44||23.37||23.30|
|2:00||21.51||21.14||22.88 (*)||23.45 (*)||23.37||23.30 (*)|
|3:00||21.51||21.14 (*)||23.82||24.42||23.37 (*)||24.58|
Table 1 - Raw Data
* indicates ignition time, blanks cells have no data
The heat of combustion of benzoic acid, as given on the label on the bottle, is 6293 calories per gram. The iron wire was also labeled with a heat of combustion of 1400 calories per gram. According to Parr, the equipment manufacturer, the energy input by the ignition module is the same during calibration and during the test run, and therefore cancels and may be ignored.
The slope of the temperature as a function of time before ignition and after thermal equilibrium after equilibrium are extremely close to zero, with variations of less than 0.01C. As a result, it can be concluded that the amount of heat transferred through the boundaries of the system was so small as to be insignificant. Random error caused by localized warm and cool areas of the bath is likely the primary reason behind temperature fluctuations, as the temperature sometimes would drop and then rise again after ignition.
From the data, the initial and final temperatures can be determined by inspection of several data points before ignition and at the end of each run. For the first run, the initial temperature was 21.51 C, and the final temperature was 23.73 C, a difference of 2.22 C. By multiplication of the amount of benzoic acid used and its heat of combustion, as well as the iron wire, the heat given off inside the bomb may be easily determined. In the case of the first run
By averaging the results of similar calculations for the first three runs, the constant of the calorimeter was found to be -2424 cal-1 C-1 (-2400 cal-1C-1 with correct significant digits).
If eq. 2 is solved for H, and the appropriate corrections for the burning wire are made, the following equation may be used to determine the heat of combustion of the unknown.
As the result of the results of this experiment are expressed in relations to grams of substance, and most tables of heats of combustion are based on molar quantities, and the conversion requires knowledge of the identity (specifically its molecular weight) of the unknown, the identity of the unknown was not determined.
With only three runs for both the standard and the unknown, it is difficult to ensure the accuracy of the experiment. However, the results for each run were in fairly close agreement, suggesting that random error was fairly small.
No substantial sources of error were identified in the course of this experiment, and as a result, no improvements were seen to be necessary.(2)
1. D. P. Shoemaker, C. W. Garland, J. W. Nibler, Experiments in Physical Chemistry, 6th ed., chapter VI, The McGraw-Hill Companies (1996).
2. R. Felder, R. Rousseau, Elementary Principles of Chemical Processes, 2nd ed., appendices, John Wiley & Sons, Inc. (1986)