If the intro paragraphs are correct, this is a great illustration of why so many subjects in math, stats and science are so much more difficult and unapproachable than they should be. Often there are no clear explanations of why fundamental principles make sense or actually work in a way normal people can understand. In this case it seems to have gone beyond that, in that there wasn't even an unclear explanation apparently.
"Surprisingly, a simple, valid explanation of the
exothermicity of combustion reactions has apparently not been
provided, [1−5] not even in textbooks on combustion [6−9] or in the
be provided in conceptual general-chemistry terms and results in a simple, predictive formula for heats of combustion.
The heats of reaction of a few specific combustion reactions have been explained in terms of bond energies, [5,7] but this has not revealed why all combustions of organic molecules are exothermic. Most bond-energy analyses have remained opaque since double bonds were treated on par with single bonds. We count double bonds as two bonds since the total number of electron-pair bonds is the same in reactants and products; only when the number of bonds remains unchanged is a general analysis per bond meaningful. On this basis, it becomes apparent that combustions are exothermic because of the unusually small bond-dissociation energy of O2 (per pair of
technical literature on heats of combustion"
I have a Ph.D. in chemistry and I have taught general chemistry at a large research university and at a small liberal arts college. For what it's worth I teach this explanation and know of other colleagues that do, and the explanation is not novel to this article. That said, it's probably less well known and less widely taught than it should be. It's counterintuitive to many people that the formation of strong bonds results in the production of heat!
> It's counterintuitive to many people that the formation of strong bonds results in the production of heat!
Not an expert in anything physics, but this strike me as a result that, while counterintuitive, should be obvious after a moment of thinking about it: as reaction tend to favor low energy states, a strong bond bond would mean low energy; and therefore the energy has to go somewhere.
Or, put another way: strong bonds require the input of more energy to break them apart, so it's not too surprising if they emit energy in the form of heat at the time of formation.
> it's probably less well known and less widely taught than it should be
This seems an important mechanism for science education content and instruction remaining wretched.
Consider "a 5-year old asks 'the Sun is a ball?! What color is the ball?!". Certainly some instructors teach it correctly. But the top 10-ish most used introductory astronomy textbooks have it wrong. And thus so do most first-tier astronomy graduate students. And this state has been stable for decades.
> the explanation is not novel to this article
Science education research is distinct from the underlying science research. If those colleagues didn't write it up and publish it, perhaps because they didn't see chemistry education research as their field... oh well.
The paper's existence makes it ever so slightly more likely some future content author gets it right, or is ever so slightly more embarrassed at having it wrong, and thus to revise it. Which over decades can sometimes move the needle. And being part of a research literature permits incremental collaborative correction, refinement, reference, and citation.
> It's counterintuitive to many people that the formation of strong bonds results in the production of heat!
A similar case. Some instructors do mention that attraction in bonds is almost all classical electrostatic attraction. Which makes this an intuitive extension of gravitational and electrostatic potential. But most instructors and textbooks don't. And so people struggle. Maybe a science education research paper or three might help. Or an interactive electron density model web app? (Something on my infinite todo list, using precomputed densities from GPAW:)
If it's a single step well, and the width of the well is fairly constant, then the steeper the slope is (stronger bonds are like stiffer springs) the deeper the well will be. You can extract the most energy with the deepest well.
It seems like gravity potential energy mental model works moderately well for chemical potentials (or more accurate enthalpy) at relatively low temperatures (vs bond energy) for statistically large numbers of molecules.
Granted it's been a few years since I studied chemistry, but isn't the word "combustion" defined as "exothermic reaction"? At least that's what wiki states. If a reaction doesn't create heat, we don't call it combustion.
Yes, in the same way an object on mercury takes more energy to reach orbit than the moon, an object falling to mercury releases more energy than it would falling to the moon.
What’s counter intuitive is what makes something a strong bond. Chemical bonds are more complex than simple gravity fields, further most chemistry involves both forming and breaking bonds. As such talking about specific bonds as strong or weak isn’t really an explanation so much as a corollary.
> This explains why fire is hot regardless of fuel composition.
This cracks me up. Despite a chemistry degree, I have never really thought to consider whether or not endothermic combustion exists, even though endothermic reactions exist. Nitrogen "combustion" with oxygen is endothermic. But really, the definition of combustion implies a high temperature, relatively rapid and self-sustaining reaction, e.g. rust and glucose metabolism aren't combustion.
If a reaction were to be endothermic would we have called it combustion in the first place ?
Implied in the name there is this notion that the reagents are looking for an excuse to combust, a release of bottled potential, this in turn implies this has to be an energy producing reaction rather than one where the reagents need to be coerced into reacting by providing external energy
> rather than one where the reagents need to be coerced into reacting by providing external energy
being endothermic or exothermic and activation energy are orthogonal.
as a matter of fact, burning wood with air is exothermic, but you don't see thees spontaneously burning until some activation energy is provided.
heck, it's not even a given that all exothermic reaction are self sustaining, there's plenty that will go on producing small quantity of heat but will required external energy to sustain the reaction; sometimes it's even the same reaction but in a different environment with a different chemistry or average temperature.
Yes, kinda, with a few small exceptions. Exothermy and being thermodynamically spontaneous are strongly correlated, but you can have endothermic spontaneous reactions, such as dissolving urea in water in those instant ice packs.
You can also drive reactions by removing the end products in lieu of providing external energy.
Right. If you increase the surface area or add more oxidizer, iron quite readily burns, see thermic lances, ferrocerium, metal powder explosions etc. But that is different than rusting.
A cutting torch for steel is really just an oxy-acetylene torch. Once the flame gets the steel hot enough, you cut back the acetylene, and the pure oxygen burns the steel away. It's a surprisingly effect method of cutting.
Wow! So the bulk of the energy we consume comes not from the hydrocarbons we're digging up, but the oxygen bonds we're breaking when we burn them. That's completely spun around my understanding of fuels. Fascinating.
That also means that the energy we use to fuel our lives mostly comes directly from recent photosynthesis, and not actually from historical sunlight embodied in the fossil fuels we're burning.
The production of heat comes from the formation of bonds, not from the breaking of bonds! Many people find this counterintuitive. The energy from combustion comes from the fact that all hydrocarbon combustion produces a large number of strong bonds in CO2 and H2O. Breaking oxygen bonds produces no energy and in fact requires energy - it's just that since oxygen bonds are weak it requires less energy than one might expect.
> Breaking oxygen bonds produces no energy and in fact requires energy
And this, in turn, is a specific example of a general principle: breaking any chemical bond always requires energy, while forming any bond always releases energy.
> It is found that anthropogenic fossil fuel combustion is the largest contributor to the current O2 deficit, which consumed 2.0 Gt/a in 1900 and has increased to 38.2 Gt/a by 2015. Under the Representative Concentration Pathways (RCPs) RCP8.5 scenario, approximately 100Gt (gigatonnes) of O2 would be removed from the atmosphere per year until 2100, and the O2 concentration will decrease from its current level of 20.946% to 20.825%.
> AFAIK our oxygen was not made recently. It accumulated over billions of years.
According to wikipedia [1], the atmosphere has 34e18 mol of oxygen, and photosynthesis produces 8800e12 mol/yr of oxygen, meaning the atmosphere turns over in a little under four thousand years. There's a separate estimate of atmospheric oxygen's "residence time" in the text, of 4500 years.
So although oxygen levels have been high for billions of years, there is a sense in which the oxygen in the atmosphere right now has only been there for thousands.
That calculation is simplistic as it is based on the current steady-state. The oxygen catastrophe was a massive change away from the previous equilibrium state which first had to deplete all the buffers. That's how we got iron ore veins, by oxidizing most of the iron out of the oceans. And there are other sinks besides iron.
Only if the reservoir is a fifo queue. :) Otherwise there's a distribution, with a tail. But oldest in atmosphere still looks (envelope scribbling...) order Myr? Rather than Gyr. Similar for breath?
> The double bond in O2 is much weaker than other double bonds or pairs of single bonds, and therefore the formation of the stronger bonds in CO2 and H2O results in the release of energy
Yeah! I remember someone talking about the methane lakes of Titan and remarking that if there was intelligent life on Titan, they would surely rule out Earth as any possible habitat for life, and would be absolutely shocked at the idea of these creatures who can only explore the rest of the universe by putting themselves inside of bags of rocket fuel, because that is what they breathe. What strange creatures those must be.
The thinking in this article also explains the logic behind the solid metal fuel -- like iron powder for example. Basically oxygen reacting with iron (rust formation) results in heat generation.
One important detail around use of hydrocarbons for fuel to power things like cars is that you need the energy to be delivered in a timely manner -- typically within milliseconds of the spark. Whereas something like a power plant could be happy to have just a hot pile of iron metal powder giving off heat for weeks at a time rather than in one massive flash explosion. Using something like iron powder that is in a pure Fe unoxidized state would require some process to get the iron into that state. Likely requiring energy -- so probably the cycle for iron powder power plants would be to use solar/hydro/wind power to generate the power liberate the oxygen from the FeO2 molecule, then transport the iron to the place where it would be used that has poor access to solar/hydro/wind, then use that there.
Anyway, articles like this really do a great job of making people think about things..
Kinda. It's the difference in entropy, mostly driven by electron affinity. In the thermite reaction, the Fe-O bond is fairly strong - rust is quite stable. But the Al-O bond is stronger. Hence once you kick it hard enough to break that FeO bond, it drives the reaction.
For air-breathing vehicles like cars, you assume the oxygen comes for free. For rocket fuel, you have to include the oxygen tanks in the energy density calculation.
You could think of energy density as the number of CO2 or H2O molecules produced by combustion per gram of fuel. If a fuel (like gasoline) can produce more H2O+CO2 per gram than ethanol, it will have a higher energy density.
No. This is explained in the article. That's only true for hydrocarbons which don't contain oxygen atoms. The energy comes from the O2 consumed. Sugar contains oxygen atoms, but those don't contribute to the energy released because burning sugar consumes less O2 per unit mass than burning hydrocarbons.
In your example, 9 moles of ethanol produces the same heat as 2 moles of isooctane, but the combustion products are substantially higher:
I don't understand your point, if you are claiming that O2 is not abundant in space, I think the idea is that we also find an ice asteroid. Where there is water, there is oxygen.
Edit:
I assume people are missing my reasoning. Yes, water requires energy to split, but assuming you are doing some other process that ends up with a resulting split anyway, you could use the hydrocarbons as a fuel source at that stage. Think recycling, not energy production.
You can produce oxygen from water, iron ore (O'Neill and other sci-fi writers somehow missed that astronauts brought back top quality iron ore from the moon,) stony rocks, etc.
To do it from water for instance you could use electricity to split up the H2O molecules; you could directly use the resulting O2 or you could use the H2 to reduce the iron ore back to H2O and thus recycle the H2. Maybe you get electricity from solar panels or a nuclear reactor.
You could burn this with cosmic hydrocarbons but you have only accomplished energy storage as opposed to an energy source. You are then competing with a sun that shines all the time and every other energy storage technology as well as nuclear fission, fusion, and decay.
The oxygen rich environment is the gift we get from plants here on Earth, ultimately they have stored a lot of "energy" in the atmosphere which we can tap.
Future wildcatters will see asteroid hydrocarbons as the material to make human bodies and cattle, everything from wood to plastics to pharmaceuticals -- every bit as much of a gold rush, but not a new energy source.
If you could make plastics from asteroid materials you could blow very large films and coat them with layers of metals and semiconductors and thus function as solar sails if not energy collectors -- these could sail to their destinations on their own power, say to be installed at the Earth-Sun L1 point to fight global warming.
Anyhow it took people a long time to get oxygen right. Cave men "mastered" fire but people couldn't make steel commercially until people understood that 80% of the atmosphere is inert stuff that goes along for the ride.
Imagine you're making a rocket engine, and both fuel and combustion products have high molecular weight. A rough example is burning CO (carbon monoxide) in O2, with reasoning being "both of them can be obtained from martian athmosphere". So the fuel isn't great, and specific impulse - Isp - is low, but if you add a relatively low-molecular weight component to fuel - like water - you can hope to have greater positive effect from lowering average molecular mass of combustion products than negative effect from adding an inert component.
That's partially illustrated by nuclear hydrogen engines - there is no chemistry, and the gas coming from the nozzle is actually colder than what conventional chemical rocket engine produces. Hydrogen doesn't participates in any transformations, it's purely a working fluid, yet Isp is about twice as good as with LOX-LH2 engines.
So at least adding water to burning process might sometimes improve the result.
Let's say you have oxygen from somewhere. You use the oxygen to burn hydrocarbons, to make energy and water and CO2.
To turn that into a closed cycle, you need to regenerate the water and CO2 into oxygen, to burn more hydrocarbons. The only way to do that is to add energy - on earth, we use solar energy via photosynthesis, but in space, you could use solar or nuclear energy via some chemical process.
But if you have the energy to regenerate the oxygen, why not just use it? Why make oxygen and then burn the hydrocarbons?
And why bother burning more hydrocarbons when a by-product of the regeneration process is carbon, a perfect fuel!
Photosynthesis is a perfect example, I know of a handful of useful chemicals that today come from the cultivation of specific plants. Food production is the other obvious one use case for it. If you're byproduct is oxygen, well, can use it for whatever you want.
A lot of carbon is made in stars and collects in dust clouds that condense into objects.
Close to the sun volatile substances such as hydrocarbons and water got cooked off so the Earth is still a dry place rich in aluminum and silicon compared to objects outside the frost line (Jupiter) which tend to have water, carbon dioxide, and hydrocarbons as major components.
Yesterday HN had a discussion about the Fermi Paradox¹. Personally, I think there is no single Great Filter, just a product of many modest filters — and one of them is that, if early life doesn't have an oxygen catastrophe², any eventual intelligent life will have a tough time not being able to burn things.
Without free oxygen, animal life could hardly develop at all, let alone intelligent life, since it would not have access to the large amounts of energy it needs to exist.
>These considerations show that atmospheric O2 stores energy originating from the sun, and that the heat of combustion in air can be regarded as fossil solar energy
This is a great article and I'm happy the author mentions "an explanation of combustion exothermicity in terms of Pauling electronegativities is not convincing". I was incorrectly taught this view and it took me some time to unlearn it.
I keep seeing comments about how certain science topics are initially presented in an overly complicated fashion. But there are two forces at play here: correctness versus accessibility. The theory presented in this paper "predicts most heats of combustion with an error of only a few percent". This is good enough for most practical applications, especially if your goal is to introduce students to this topic.
But this is not the most "correct" description of reality that we currently have. A better model to decrease the error would incorporate information about the 3D structures/conformations of the molecules and some funky terms related to the quantum mechanics.
Anybody publishing literature about/teaching something like organic chemistry is stuck between this balancing act of correctness versus accessibility. This difficulty is compounded by the fact that you basically have to unlearn some of the stuff you picked up during the path to "accessibility" in order to properly move into the "correctness" phase. I genuinely sympathize with anybody dealing with the balancing act.
After studying organic chemistry for an extended period of time, it dawned on me that the well-thought-out explanations in my textbooks were just post-hoc rationalizations the field uses to avoid delving into the true quantum mechanical nature of the reactions. I'm happy I sacrificed some correctness for the huge amount of accessibility I got. But I'm also happy to unlearn some of the stuff on my path to correctness.
I feel like this explanation is a good step but misses the crux that any combustion process (even with a non-oxygen oxidizer, if that's still called combustion) exists because each broken bond releases enough energy to break the bond of neighboring atoms/molecules with some left over to convert to heat.
The heat is a byproduct of the cascade reaction that allows combustion to occur, which structurally in my mind is very similar to fission processes where neutron bombardment produces excess neutrons to bombard more atoms.
"Surprisingly, a simple, valid explanation of the exothermicity of combustion reactions has apparently not been provided, [1−5] not even in textbooks on combustion [6−9] or in the be provided in conceptual general-chemistry terms and results in a simple, predictive formula for heats of combustion. The heats of reaction of a few specific combustion reactions have been explained in terms of bond energies, [5,7] but this has not revealed why all combustions of organic molecules are exothermic. Most bond-energy analyses have remained opaque since double bonds were treated on par with single bonds. We count double bonds as two bonds since the total number of electron-pair bonds is the same in reactants and products; only when the number of bonds remains unchanged is a general analysis per bond meaningful. On this basis, it becomes apparent that combustions are exothermic because of the unusually small bond-dissociation energy of O2 (per pair of technical literature on heats of combustion"