LAD 10.3 (pg 1 of 2)
Enthalpy of a Combustion Reaction
Introduction
In this lab, the enthalpy of the combustion, ∆Hcomb, of butane will be measured experimentally and calculated theoretically. This is
distinctly different from specific heat capacity. It is important to realize that heat emanates from the butane because it is reacting, not
because it is high in temperature and cooling down. The heat is produced by burning butane, and this means that we need a different
equation to calculate heat, q.
The equation is q = (∆ H )(amount) .
∆H is the symbol for enthalpy, which is the heat (gained or lost) for a process in
which the product of the pressure and volume of the system is held constant.
The units on ∆H are heat/amount. The amount is usually moles, but it could be other amount units such
as grams or kilograms.
In this lab we need to catch the energy that is evolved while butane burns and this will be done by a known volume of water. Again we
will make the (not so good) assumption that the heat lost by the butane will be gained by the water.
PRELAD
Set up a data/results table. Be sure and make a row for each process the data item.
Procedure
1.
2.
3.
4.
5.
6.
First light your portable butane burner to be sure it works and determine the correct height for your can set-up.
Shut off the burner and measure the mass of the portable bunsen burner. (If the mass is more than the balances capacity, you may
need to unscrew the burner unit on top. If you do this, be sure and mass if afterwards without the burner as well.)
Tare the can and measure the mass of a half a can of cool water from the tap.
Measure the starting temp of the cool water in the can.
Heat the water (up to some temp above 60ºC and below 90ºC) by burning the butane. Immediately after shutting off the burner,
record the hottest temp of the water.
After shutting off the burner, measure the mass of the portable burner after completing the burn.
Processing the data
a.
b.
c.
d.
The water is simply changed temperature. Calculate the heat gained by the water.
Do a simple subtraction to calculate the mass of butane burned.
Use the molar mass of butane to convert the mass of butane burned to moles of butane burned.
The water changed temperature because of the burning butane, thus we will assume that the heat gained by water is equal to the
heat produced by butane. Thus in part a) you actually calculated the heat given off by burning the amount of butane that you
burned. Use that number with the mole values from part c) to calculate the energy evolved per 1 mole of butane burned. This is
the experimental value of ∆Hcomb.
Post LAD Questions
1.
The “Process the data” section took you step by step through the calculations. Alternatively, you could have done a literal solve of
the the following equations and simply substituted your measured data.
Formulas that you know and love:
A.
moles * molar mass = mass
B.
qwater = cw * massw * ∆Tw
C.
qbutane = ∆H * moles
Set equation B and C equal to each other. Then solve equation A for moles and substitute into the combined equation you just
wrote, then solve this “literal equation” for ∆Hcomb. (This means solve it for ∆H using just the variables.)
LAD 10.3 (pg 2 of 2)
Enthalpy of a Combustion Reaction
2.
Is burning butane and exothermic or endothermic reaction?
Scientists have agreed that exothermic reactions will be symbolized with negative ∆H values and endothermic reactions will be
symbolized with a positive ∆H values.
After the calculations should we apply a + or – sign to the molar enthalpy that was calculated in “d”.
3.
Write a balanced equation for the combustion of butane.
4.
Now we need to determine a theoretical value for the burning of butane.
We can go to the literature and look up a heat of formation value, ∆Hf for any substance listed.
∆Hf defined:
The energy (lost or gained) for the reaction that represents the
formation of that particular substance from its elements.
Using the ∆Hf values listed on your Thermo Table, we can apply the following formula.
∆Hrx = ∑n∆Hf products − ∑n∆Hf reactants
• ∑ means calculate the sum (add together)
• n stands for the stoichiometric coefficient of the
substance in the balanced equation.
• ∆Hf stands for the enthalpy value looked up on the list
for formation reactions for the particular substance
Don’t forget that ∆H values are extensive, and that they are reported per mole of some molecule. You must be sure and multiply
the ∆H values times the coefficient, n (from the balanced equation) of that substance.
5.
The theoretical value for the combustion of butane is 2661 kJ/mol. Calculate a percent yield for the molar combustion of butane.
We will take this % yield as a measure of the % of efficiency of our ability to utilize the chemical potential energy of the butane.
6.
How is the precision of the class data? How is the accuracy? Does it appear if the results are exhibiting random or systematic
error? Explain.