A worked example went something like this:

Use the following experimental data to determine the enthalpy change when $\pu{1 mol}$ of ethanol ($\ce{C2H5OH}$) is burnt:

Mass of water = $\pu{150.00 g}$

Initial temperature of water = $\pu{19.5 °C}$

Maximum temperature of water = $\pu{45.7 °C}$

Initial mass of spirit burner = $\pu{121.67 g}$

Final mass of spirit burner = $\pu{120.62 g}$

Being a worked example, I saw the worked solution and understood the process completely. What's bothering me is a fundamental flaw in the procedure that has seemed to be neglected completely.

We are relying on the idea that the combustion of ethanol will produce some energy, the amount of which can be calculated with the increase in temperature of the beaker of water on top. However, if there was no ethanol, no combustion, the flame giving rise to the combustion would still manage to heat the water up; in this sense, we do not know whether it is the flame supplying the energy to the rise in water, or the exothermic reaction that takes place with ethanol. How does the existence of such a flaw make the procedure useful to any extent? Is there anything I'm missing?


2 Answers 2


I see your concern, and it's absolutely valid. Any time you perform a calorimetry experiment, you need to be certain that you're not introducing extra energy into the system. Typically, the spark used to ignite the sample has a negligible energy input, but it is still there. However, any heat released by the flame (other than the initial energy put in to generate the spark) is generated by the combustion process. Without the combustion, the flame would cease to exist. If you were running this experiment, you could use an electric sparker and measure the power input for the sparker to know how much energy was inputted. Another option would be to run a control experiment with just an ignition (but no ignitable material) and see how much, if any, the water heats up.

Of course, if you used a lighter to start the flame, and then held the lighter open throughout the experiment, you would be introducing extra energy from the combustion of lighter fluid, and this would give you false results as you feared. That's why we don't run calorimetry experiments in that way.


Change in heat=mass of water × specific heat capacity × change in temperature But; mass of water=150•00g Specific heat capacity=4•2 Change in temperature= final temp - initial temp =45•7 - 19•5=26•2 Change in heat=150×4•2×26•2 =16758 But mass of ethanol burnt = initial mass of spirit bottle - final mass of spirit bottle =121•67-120•62 =1•05 Relative molecular mass of ethanol(c²H⁵OH)=((12×2)+(5×1)+(16×1)+(1×1)) =46 Moles of ethanol burnt=mass burnt÷relative molecular mass =1•05÷46 =0.022826087 0.022826087moles of ethanol produces 16758 joules I mole= (16758÷0.022826087) ×1 =734,159.9986016 •°•, 1 mole produces approximately 734,160 joules per mole 734,160 joules per mole =734,160 joules per mole÷1000=734•160 Kilojoule per mole C²H⁵OH + O²-> 2CO² + 3H²O; Change in heat = 734•160 KJ/mole

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