Ethanol in Gasoline, a Technical Perspective
I originally wrote the following for a message board back in 2015. My thinking has not changed.
There seems to be a universal hatred for the use of ethanol in gasoline on this board. In my opinion, a little ethanol in gasoline is not a bad thing. If you are not interested in 15,000 words on why and just want your bike to run better, skip ahead to the section titled Corrective Action. I plan to steer clear of the political and environmental aspects of ethanol and focus on the technical ones (although economics will undoubtedly enter the picture). My thinking is that since E10 is gasoline where I live and its ethanol content is already subsidized, I may as well make the best of it.
Back in the early days of aviation when the maximum compression ratio on gasoline was around 5:1, ethanol was recognized as a “superior” fuel because it allowed a compression ratio of 7:1. But it was comparatively expensive and required a much greater volume to do the same work (in other words, the miles per gallon was poor). Long before the idea of octane rating was even conceived, ethanol was known to inhibit engine “knock.” Today we assign pure ethanol an octane number of 98 based on the (RON+MON)/2 method (which is the method by which pump gas is labeled in the USA). However, its Blending Octane Number is much higher. When ethanol is added to gasoline, the resulting octane rating of the blend is more than a straight linear blending equation would predict. For example, 10% ethanol in gasoline typically yields a three-number increase in RON and a two-number increase in MON. This makes ethanol a very economical octane booster and – in my opinion – offers a compelling argument for its inclusion in gasoline.
Probably the best layman's description I have heard for gasoline is that it is a “chemical soup.” Pure gasoline (without ethanol) is a mixture of hundreds of different hydrocarbon compounds. Hydrocarbons contain hydrogen and carbon atoms in different quantities and arrangements. They can vary from quite light with 4 carbon atoms to fairly heavy with 12 carbon atoms. An example of a light hydrocarbon molecule is the gas butane (which is commonly added to gasoline in the winter months). Butane has a chemical formula of C4H10. For thermochemical calculations, a useful approximation of gasoline as a single hydrocarbon compound is C8H15.
In contrast, pure ethanol, which is an alcohol (and exactly the same stuff you drink) is a single substance. It is also made up of hydrogen and carbon (so it is a fuel), but contains oxygen too (this is the reason it is referred to as an “oxygenate” when added to gasoline). Its chemical formula is usually written as C2H5OH. Ethanol is not the only oxygenate to have been used in gasoline. Over the years, ethers like TAME (C6H8O), MTBE (C5H12O), and ETBE (C5H12O) have also been used.
A quick digression about beverage alcohol versus motor fuel. A motor fuel must be “denatured” (rendered unfit for human consumption) before leaving the manufacturer. Denaturing is accomplished by adding a small amount (maybe 1 - 2%) of gasoline to the ethanol before it leaves the ethanol plant. This is entirely a tax thing as beverage alcohol is taxed at a higher rate than motor fuel.
Another alcohol you may recognize is methanol (aka wood alcohol) – chemical formula CH3OH. This fuel has long been used in racing. It is famous for its intake-charge cooling effect (due to a high latent heat of evaporation). Engines drinking methanol can run comparatively high compression ratios. I wrote “drinking” because it is also famous for high fuel consumption – more than twice that of gasoline.
Hydrogen and carbon are the two elements most commonly used as fuels. Sulfur could be used as a fuel too, but its drawbacks far outweigh any benefits. In fact, chemical engineers go to great lengths to exclude sulfur from hydrocarbon fuels because of the corrosion damage it can do.
When hydrogen and/or carbon combine with oxygen from the atmosphere an exothermic (heat-releasing) chemical reaction occurs. The reaction is said to be stoichiometric (from Greek, meaning “first principle measure”) when all the fuel and all the oxidizer combine completely. You can think of it as the ratio at which the hydrogen, carbon, and oxygen atoms all find dance partners. Ideally, only water vapor (H2O) and carbon dioxide (CO2) are released. None of the original fuel or oxidizer remains unreacted. By the way, this never actually happens inside any real engine. In a gasoline engine, as much as 1% of the oxygen never gets to combine with fuel no matter how rich you make the mixture. Air is about one-quarter (23.2%) oxygen by mass. So something like 4% of the air inducted can go unreacted – but stoichiometry is a nice expedient.
The stoichiometric air-fuel ratio (AFR) is about 14.6:1 by mass for gasoline. I say about because it depends on the exact hydrocarbons present in the gasoline. In other words, 14.6 pounds (or grams, or any mass unit you like) of air is consumed per unit mass of fuel. Whereas for ethanol, the ratio is 9.0:1, exactly. Remember, ethanol is a single substance rather than a chemical soup. An important point to note is that generally speaking, the smaller the AFR number, the “thirstier” the engine will be when run on that fuel – the lower the miles per gallon (MPG) it will get.
It is also important to note that a stoichiometric mixture is often discussed solely because it is chemically easy to analyze. (It is also the only mixture ratio where a 3-way catalytic converter can clean up exhaust emissions, and this is the real reason it is used in the automotive world.) Maximum torque is always produced on the rich side of stoichiometric (a smaller AFR number). People usually want to see a number they can relate to, like 12.5:1. I prefer to stay away from that and just say “richer” because the exact AFR for maximum torque varies according to a lot of factors – not the least of which is how it is being measured. Similarly, the best fuel economy is found on the lean side of stoichiometric because that is where the highest thermal efficiency occurs. (You get the biggest push on the piston per unit of fuel burned.)
You may question my use of the word “torque” throughout this write-up. If you are uncomfortable with that, replace it with horsepower. I started out by writing horsepower (which seems to be more readily understood), but it is also one step further removed from what is actually changing. Power is the time-rate application of torque. Power simply equals torque times RPM (with an appropriate scaling constant employed, depending on the units used).
The torque produced by an engine is determined by a complex array of factors, but the single most important one is how much air it can ingest (and retain). The more air in the cylinder, the more fuel that can be burned, which releases more heat, which causes more pressure, which yields a more forceful push on the piston.
An oxygen-bearing fuel (like ethanol) can be thought of as a “chemical supercharger.” That is, the fuel itself brings some oxygen to augment the oxygen in the air. The most dramatic example of this concept is the use of nitromethane (chemical formula CH3NO2) in top fuel dragsters. The stoichiometric AFR for nitromethane is 1.7:1. Talk about poor MPG!
Speaking of chemical superchargers, another example of such is nitrous oxide (N2O). Returning to the world of piston-engine aircraft (incidentally, this is where all the really good pioneering research on internal combustion engines occurred – automotive and motorsports applications were just trickle-down applications) where attempts were made to introduce pure oxygen into the intake plumbing. The result was extremely high combustion temperatures that quickly melted pistons and valves. It turns out that all that “inert” nitrogen in the air is necessary. Although nitrogen plays no part in the combustion reaction (other than forming some oxides of nitrogen at very high temperatures), it does add mass to the “working fluid” inside the cylinder. This mass gets heated along with the combustion products. Its expansion pushes on the piston while helping maintain a reasonable temperature. Thus, with nitrous oxide, there are two atoms of temperature-moderating nitrogen “along for the ride” with each atom of oxygen.
The chemical combination of (and equation for) fuel and oxidizer is by mass, but liquid fuels are sold by volume. So, just knowing the stoichiometric ratio is insufficient to calculate fuel consumption. Before I get into the arithmetic, note two things:
Air-fuel ratio is typically a “gravimetric” (based on weight) measurement, but it is equally valid to talk about it in volumetric terms. In fact, this will make comparisons of different fuels easier to follow.
The reciprocal of air-fuel ratio (AFR) is fuel-air ratio (FAR). If the air-fuel ratio is 10:1, the fuel-air ratio is 0.1:1 (0.1 unit of fuel per unit of air).
In order to see how much fuel gets burned, we need to do a fuel volumetric analysis. For that, we need to know the density (aka specific gravity) of the fuel. This is easy to determine for a pure substance like ethanol (0.79) or nitromethane (1.139). The specific gravity of gasoline is a little more variable (remember it is a chemical soup), but 0.74 is a reasonable value. See the table below.