The heat of combustion is the energy released when a certain amount of material, usually one mole, reacts exothermically with oxygen in the air. It is an essential thermochemical property in the characterization of fuels, alcohols and, especially, hydrocarbons, such as butane and methane.
When a substance burns, it releases light and heat, energy that can be harnessed to do work on other bodies. For example, the heat generated is capable of heating large volumes of water, the steam of which presses on the surroundings, driving electrical generators; or simply by heating the steamed food even more.
The heat of combustion varies between all substances, even though their chemical nature is essentially the same. This variation corresponds to the relative stabilities, that is: the more unstable a compound is, the greater the energy released, represented as -ΔH. The negative symbol means that the heat flows out.
Different values of -ΔH are used to compare the relative stabilities between different alkanes and their isomers. Likewise, it shows how stable an oxygenated compound, such as a ketone, is against another, an aldehyde or carboxylic acid .
To measure the heats of combustion, a calorimetric bomb is needed. Inside it, the substance reacts with oxygen by activating an electrical spark.
The released heat then heats a volume of water surrounding the sample compartment, measuring the temperature before and after combustion, as well as the mass of water and fuel.
So, the released energy will be equal to:
C EH2O · m H 2 O·ΔT
Where C eH2o is the specific heat of water, 4.184 J / g · ºC, m H 2 O the mass of water, and ΔT its change in temperature. Finally, this heat, expressed in units of joule or calories, is divided by the mass or moles of the fuel placed inside the calorimeter pump, to obtain the heat of combustion per unit of mass or of moles.
Metals do not burn themselves, but oxidize at different rates depending on the temperature to which they are exposed. The heat they produce is negligible to be measured in the same way as is done with hydrocarbons and other fuels. They are therefore not combustible substances.
During combustion the bonds break to form new, more stable ones. The energy contained in all the bonds of a molecule is independent of its movement, which is why we speak of a potential energy.
To know exactly what these new bonds are, one must consider the products of complete combustion: CO 2 for carbon, H 2 O for hydrogen, and NO 2 for nitrogen. As regards hydrocarbons, their complete combustion will generate stoichiometric mixtures of CO 2 and H 2 O.
Both molecules, CO 2 and H 2 O, have very low potential energies, because their bonds (O = C = O and HOH) are very stable compared to the bonds of the hydrocarbons from which they came (CH and CC). .
The potential energies are immeasurable. But their variations are not, that is, the difference in these energies between the products (CO 2 and H 2 O) and the reactants (hydrocarbons).
That is why having these variations, or what is the same, their heats of combustion, it is possible to know what the relative stabilities are between a set of hydrocarbons or isomers.
Alkanes can be linear, branched, or cyclic. In the case of linear alkanes, their heat of combustion varies as a function of the length of their chains; that is, it depends on how many CH 2 units they have. Consider the example of n -hexane, n -heptane, and n -octane:
CH 3 (CH 2 ) 4 CH 3 , -ΔH = 4163 kJ / mol
CH 3 (CH 2 ) 5 CH 3 , -ΔH = 4817 kJ / mol
CH 3 (CH 2 ) 6 CH 3 , -ΔH = 5471 kJ / mol
It can be seen that their heats of combustion vary by 654 kJ / mol. This means that each CH 2 added to the chain increases the heat of combustion at a rate of 654 kJ / mol. Thus, nonane, CH 3 (CH 2 ) 7 CH 3 , will have a heat of combustion equal to 6125 kJ / mol (5471 kJ / mol + 654 kJ / mol).
This is the same as saying that longer chains have higher potential energies, and therefore are more unstable.
Branched alkanes are more stable than linear alkanes, and this is deduced from their heats of combustion. Now consider the case of three isomers of octane:
The n- octane is unstable because its heat of combustion is the largest (5471 kJ / mol). In contrast, the 2-methylheptane and 2,2-dimethylhexane isomers have lower heats of combustion (5466 kJ / mol and 5458 kJ / mol, respectively) because they are branched. 2,2-dimethylhexane is the most stable because it is the most branched of the three isomers.
The relative stabilities of alkenes can also be obtained from their heats of combustion. Consider, for example, four isomers of butene:
Here we no longer speak of ramifications, but of the degree of substitution of the double bond, C = C, and its van der Waals tensions; that is, of the electronic and steric repulsions between two bulky groups located on the same side of the double bond.
The first isomer, 1-butene, is the most unstable because its double bond is less substituted (H 2 C = C). This is reflected in its heat of combustion of 2717 kJ / mol.
On the right, cis-2-butene is a bit more stable because its double bond is more substituted. But trans-2-butene is even more so, because although it has the same degree of substitution, the CH 3 groups of the double bond are in trans positions, that is, on different sides of the double bond. Note that the heat of trans-2-butene (2707 kJ / mol) is only slightly less than that of cis-2-butene (2710 kJ / mol).
The last isomer, 2-methyl-2-butene, which is also just as substituted as cis and trans 2 butene, is however the most unstable of all. This is because both CH 3 are found on one of the carbons of the double bond, thus being an alkene with geminal CH 3 .
As with alkanes, alkenes, and other hydrocarbons in general, the relative stabilities of carbonyl compounds (ketones, aldehydes, carboxylic acids) are also related to their heats of combustion.
Thus, the more stabilized, or less reactive, its C = O groups, the lower its heats of combustion.
For example, the heat of combustion of butanal, CH 3 CH 2 CH 2 CHO, is 2475 kJ / mol. This heat is higher than that of 2-butanone, CH 3 CH 2 COCH 3 , which is equal to 2442 kJ / mol. Therefore, butanal is more unstable than 2-butanone.
In 2-butanone, C = O is more stabilized thanks to the CH 3 CH 2 and CH 3 groups ; This is not the case with butanal, where one of the groups is barely an H atom.