Experiments with
Butane L i g h t e r
Fluid
The availability of disposable butane lighters and of 4-oz refill eannisters for non-disposable butane lighters make possible three useful semi-quantitative experiments (or demonstrations) for introductory chemistry courses. These involve the determination of average molecular weight, a study of the temperature variation of vapor pressure using a tire gauge and, least accurately, the molar heat of combustion. All our data were obtained with "Scriptom Premium Butane Fuel" and "Crickete" disposable lighters but other brands should serve equally well. Average Molecular Weight of Butane Llghter Fluid The 4-oz refill cannister is weighed on a standard analytical balance. Tygon tubing is then used to connect the valve (without adapter) to a bent piece of glass tubing which can deliver vapor t o a graduated cylinder arranged for downward disnlaeement of water. The valve is eentlv deoressed until a convenient volume of vanor has been collected. The deliverv tube is then carefully disronnected and ;he ontainer reweighed. Calculated averagi molecular weights fall in the ran& 57 i 2 amu. The volume of the butane vapor should be read as RMn as temperature equilihrlum has hprn reached sinre the equilibrium solubilny of hutane in water is surprisingly large. The kinetics of dioiulurian, howwrr. m e fortunntrly inordinately slow and the results are seen t o err rather on the low than on the high side ~
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Use of allre Gauge to .Measure the Vapor Pressure ot Butane Lighter F b l d as a Functlon of Temperature By means of a hack-saw and a small round file one of the adapter nozzles supplied with the refill cannister is made t o fit snugly hoth over the valve and inside a standard tire gauge. The pressure inside the cannister is then measured in the normal way a t various temperatures hetween 0' and 50'C and the gauge pressures in psi's are converted t o absolute atmospheres. The metal container facilitates heat transfer with the various temperature haths, hut care must he taken t o allow sufficient time both before and between readings t o ensure thermal equilibrium. In the table values obtained Vapor Pressure of Butane Lighter Fluid as a Function of Temperature Literature (atm) interpolatedb ~s~~u~ateda
Temperature" ( K )
Gauge Reading (psi)
A b ~ o l v t ePressure ( a m )
2.4 14.4 26.0 35.0 40.5 45.0 50.0
12 23 37 54 62 70 81
1.8
1.68
2.6 3.5 4.7 5.2 5.7 6.5
2.50 3.54 4.54 5.25 5.88 6.65
1.7 2.5 3.5 4.5 5.1 5.8 6.6
acalculatea from formula given in: "API Selected Valuer of Propertier of Hydrocarbonr." b InterPolatea from graph given in: JoLdan. T. Earl. "Vapor Prerrure of Organic Compounar". Wiley-Interscience. YorK. 1954.
New
usine both an automobile tire eauee (0-30DC1and a Sebwinn hicvcle tire eaure are cornoared with literature .. .. (35-55%) . valuk for iso-hutane It war tGe gauge readings which first told &(what we should have'knuwn anyway) that 'butane" lighter flwd is principally i.w- hutane. A glc analysis af two samples gave: 9.100 ,so- hutane. 5% n -butane and 1% (prohahly) propane. It is, however, unlikely that identilicetion of isomers a i t h a tire gauge wdl bwumr standard methMolar Heat of Combustion of Butane Lighter Fluid Disousahle butane lighters are light enuurh to be weiehed on an analvtical balance. The heat of combust;on of irobutane i n be approximated by allowing the flame t o raisethe temperature of a known amount of water. The simple apparatus used for finding the heat of comhustion of a candle in the CHEM Study Laboratory Manual is adequate for this admittedly rough determination. Adjustment of flame height and variation of experimental arrangement can cause the experimental values of the molar heat of combustion to vary over a considerable range. Even under the best conditions substantial heat loss occurs and total comhustion palpably does not take place since some carbon black is still deposited. An average experimental value of 440 kcallmole is to be compared with the literature i-Cdho(g)
+ 1312 odd
-
4COzk)
+ 5HzOk)
AH = -628.34 kcal
Even this lower limit value for the molar heat of comhustion of isa-hutane is, however, nearly 100 times larger than the molar heat of vaporization calculated from a Clausius-Clapeyron plot of the vapor pressure data given ahove. Such order-of-magnitude comnarisons of heats relatine " - to the hreakine of van der Waal's "bonds" and those relatine to the making and breaking of chemical bonds are in same ways more pedogagically important than very precise values of either. At the very least they might have prevented the authors of a recent teat from giving a value of 539.55 kcal mole-' far the molar heat of vaporization of water! Safetv Note ~~~Some recently purchased cans of lighter fluid have been found to have vapor pressure considerably higher than that of hoth normal- and iso-butanes, no doubt due to the presence of very small amounts of propane and even of ethane and methane. The observed molecular weight obtained from these hiph-pressure cans, however, remained within the limits mentioned above as a full can was pro&essively emptied. The observed values of the vapor pressure fell progressively to the expected values. A second "high-pressure" can was heated until the Lapped-metal seams burst. This occurred a t 9O0C. There would seem no danger, therefore, in immersing cans a t 50°C. Derek A. Davenport P u r d u e University West Lafayette, Indiana 47907
306 I Journal of Chemical Educaflon
