Heating values of fuels: An introductory experiment - Journal of

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Heating Values of Fuels An Introductory Experiment Timothy R. Rettich Illinois Wesleyan University, Bloomington. IL 61702 Rubin Banino' and David J. Karl2 Wright State University, Dayton, OH 45435 T h i s experiment is a simple, inexpensive way for students t o determine t h e heats of combustion of common solid, liquid, and gaseous fuels. T h i s straightforward calorimetric procedure has been used successfully i n our introductory courses a n d is versatile enough to be a n appropriate experiment in physical chemistry. We have examined all hack issuesof this J o u r n a l a n d found n o comparable experiment. Although dated, readers may be interested in the survey of thermuchemical experiments by Neidig et al.?

The Experiment The calorimeter (Fig. 1) is assembled from a 250-mL Erlenmeyer flask containing 200 g of water and a a 0 'C thermometer, which also serves as stirrer. To prevent breakage a small piece of rubber tubing fits over the tip of the thermometer bulb. As shown in Figure 1, the flask is shielded from drafts by a chimney made from a l-lb. coffee can or any can of similar size. A few notches cut into the circular hole (chassis punches can be used for this) in the can channel the air flow. For reproducibility, the bumseontinuelong enough to produce a 20 K rise in the temperature of the water. Gaseous, liquid, or solid fuels may be used as described below. The calorimetric system is calibrated by burning a fuel with a known heat of combustion. We recommend the use of methanol in an alcohol bumer (such as Fisher Scientific's #04-245-AA). With the burner about one-third full and with its cap on, it is weighed to

Buret (50 mL)


Rubber Bulb Soap Soln.

Figure 2. Flowmeter f w g a ~ r n ufuels. ~

f 0.01 g both before and after the burn. This yields the mass of fuel consumed and permits the calculation of the heat loss factor for the system. The alcohol burners can be filled with many different liquid fuels-see the table. The weight loss of a solid fuel can he directly measured. A candle supported on a watch glass can he weighed before and after burning. Different types of wood can he used as fuel; our hest data have resulted from matchsticks (without heads) built into a five- or sixlevel "campfire" on top of a block of solid wood. I t is helpful to light the campfire at several points along the perimeter and let it burn toward the middle. The residue can then be reweighed. The amount of gaseous fuel used can be determined from the time of bum. the volume flow rate. ambient temoerature and pressure, and thiidesl eas ..~law. A simoie flowmeter td measure theiuel flow rare through the burner (Fig. 2) can beeasily cnnstructed. A burner without air inlets at the haxe is required here, or the upenings in a convenrional hurner ran be realed. Care must he taken lo wurk at moderate and reproducible flow rates. The mass of the dry flask and the mass ofwater are determined by weighing to f0.01 g. Some practice is requried in lighting andextinguishing flames so as to be reproducible. Stirring with the protected thermometer should be continuous during the burn. The final temperature, measured to f0.1 K, is taken as the maximum tempera~~~~~





' To whom correspondence should be addressed.

Deceased. Neidig, H. A,; Schneider, H.; Teates, T. G. J. Chem. Educ. 1965,

42. 26.

Figure 1. Calwimeter fw determining heats of combustion of fuels. Alwhd burner can be replaced by other combustion sources. 556

. I n ~ m a of l Chemical Education


Miner, D.: Seastone, J.. Eds. Handbmk of Engineering Materials; Wiley: New York, 1955.

Heats of Combusllon of Common Fuels (kJ mol-' unless otherwise noted)

Methanol Ethanol 1Propnol l-Butan01 ~Hexane rrHeptane 22.4-Trimethyl Pentane Gasoline Yellow Vdlve Candle Whit* Emergency Candle Wwd ~~

1303 1919 2749 3047 4163 4961

2.5% 4.5% 11% 35% 20% 13%

4.7% 5.0% 2.7% 27% 13% 9.1

1.3% 2.4% 2.0% 2.4% 0.4% 1.9%

0.7% 2.9% 0.7% 0.9% 7.3% 17%


726.5 1366.8 2019.8 2675.8 4163.1 4811.2 5455.5


50.06a 14.20d

Methane Ethane Propane *Butane

5.8% 17%





2.1 %

23 %





4.6% 2.0% 2.3% 2.1%

2.7% 3.3% 8.3% 2.5%

18.4ge 890.31 1559.8 2219.9 2877.0


ture. This is usually attained within 30 s of extinguishing the flame. Use caution in dealing with the flammable materials and hot objects used in this experiment.


The heat loss factor f of the calorimeter can he calculated by heat generated from fuel ( Q J (1) = heat absorbed by apparatus (Q,) where Q1, which is for 1: moles, is determined from literature values in kJfmol for the standard fuel (e.g., methanol) and the amount of fuel burned. The heat absorbed by the flask and water, Q2,is determined by Q, = (mass of flask X sp. ht. flask X

+ mass H20

sp. ht. H20)X ( A T calorimeter) (2)

Note that the specific heat of Pyrex is 0.858 J g-' K-'.4 In -~~ tests with all other fuels. the heat absorbed,. Qs, as determined using eq 2, is multiplied by the heat loss factor/ to determine the heat eenemted from the fuel. The value off, typically between 1.3and 1.7, should be constant for a given a o ~ a r a t u ssince rouehlv the same quantity of heat is produked during every test: The value of H in standard units of heat per quantity of fuel consumed, n, can be calculated by ~~~

1376 1961 2656 4126 4459 4532


AH = V X Q,)/n For solids and liquids, n is easily calculated from the mass and molecular weight. For gaseous fuels, n can he calculated using the ideal gas law, where P and T a r e ambient pressure and temperature, and the volume of gas used is the product of the time of burn and the volume flow rate. (Note that the enthalpy change is the negative of the heat of combustion as defined above.)

866.1 1508 2036 2805


Results and Dlmussion

The average results for a variety of fuels are shown in the table. Column A represents a compilation of student results, column B are the authors' results, and column C lists literature values. Careful work allows an imprecision of less than 570, and beginning students can usually achieve 10%. The accuracy is far more variable, and depends on the fuel used. There is a difficulty with soot formation as more complicated hvdrocarhons. such as easoline. are burned. Soot formationcan be lessened by azjustingdhe wick height, but this involves lenethenine the time of burn. which then adverselv influences t i e constancy of f. For introductorv courses, this ex~erimentallows the student to measure several types of data and use them to calculate "real" results that can he compared to literature values. The simplicity and economy of dksign for this experiment allow its use in large lab sections. This experiment also served as a startin; ooint for class discussion of several issues, including the relation between heating values and the chemical comoosition of fuels. the comoosition of easoline. and what is mkant by "octanenumber": When used in a ohvsical chemistrvlahoratorv. this e x ~ e r i ment proved to b; valuable lesson in error analysis.- he lareer number of fuels that could be examined bv advanced stubents in one period also allowed more detailid chemical comparisons to be made, such as examining homologous series. This direct experimental determination of enthalpy changeof combustion is inmany ways agood preparation for and a useful supplement to the classical bomb calorimeter experiment. We should note in passing that we tried many versions of the basic design to minimize the f factor or to increase the reproducibility, and ended up with our original design (Fig. 1). Students should be encouraged to suggest andlor test other designs.





David J. Karl unexpectedly died during the preparation of this manuscript. We owe him a profound debt of gratitude for his encouragement, enthusiasm, and skill.

Volume 65

Number 6

June 1988