When did fire become a thing? No one knows the answer to that question. Fusion certainly occurred before fire—it happens in suns, along with nuclear fission (radioisotopes exist in the sun)—but this is not fire. It appears flamey. It is hot. It radiates through varying segments of the electromagnetic spectrum. But I am going to limit the definition of “fire” to “combustion,” if you don’t mind.
The simplest combustion reaction occurs when pure hydrogen (H2(g)) and oxygen (O2(g)) gasses are combined in a 2-to-1 ratio and given a little energetic push called activation energy (i.e. hydrogen and oxygen will hang out with each other unless they are provided this energy). Diagrammatically, the activation energy looks like this:
The reactants (hydrogen and oxygen in our example) start on the left side of the hump, an appropriate (or excess) amount of energy is provided, and products result on the right side of the hump. The “ΔH” thing on the right side is beyond the scope here but represents a positive, negative, or neutral amount of energy released in the reaction.
The amount of activation energy varies widely from very small (e.g. some explosives) to “no reaction will ever happen regardless of energy input.” Here is what the most basic combustion reaction looks like in chemical reaction shorthand called “stoichiometry:”
2H2(g) + O2(g) → 2H2O(g)
And now, an entertainment of limited scientific value:
Combustion is generally thought to involve hydrocarbons (e.g. octane in the “gasoline” or “petrol” you use in automobiles) or their oxygenated friends the carbohydrates (e.g. cellulose, a polymeric carbohydrate used in paper and present in wood). The simplest combustion reaction is between methane (CH4(g)) and oxygen (2(g)), again resulting water but also resulting in carbon dioxide (CO2(g)) when the reaction occurs efficiently. When it does not occur efficiently or when it occurs in the presence of other substances (e.g. most of the time) it produces by-products including carbon (elemental symbol “C” aka “soot”). Here is the stoichiometry of that simple reaction:
Methane is commonly known as natural gas, although natural gas is not pure methane when used as a fuel. What the stoichiometry tells us about this reaction is that each molecule of methane uses two molecules of oxygen and produces one molecule of carbon dioxide and two molecules of water, along with an amount of energy released in the process. The energy is used to heat various processes, including home furnaces and water heaters, and used to drive steam and gas turbines to produce electricity.
When octane is used as the hydrocarbon, the balanced equation is as follows:
2C8H18(g) + 25O2(g) → 16CO2(g) + 18H2O(g)
In common English, this means that each molecule of octane requires 25 molecules of oxygen (and that activation energy thing, typically supplied by spark plugs) and results in 16 molecules of carbon dioxide and 18 molecules of water, along with a good burst of energy that drives the pistons, drive shaft, and wheels; the wheels have tires that turn and exert a force against driveways, roads, dirt, mud, water, etc. and the automobile moves forward—or backward—at various speeds as allowed by the transmission.
Candles (if you were wondering where all this leads) are made from paraffin wax, which is a varying mixture of hydrocarbons typically with between twenty (C20) and forty (C40) carbons in their structures. A C20 hydrocarbon like eicosane can have up to 366,319 isomers (isomers all have the same chemical formula of a chemical compound but differ in physical and some chemical properties), while tetracontane (C40H82) has 62,491,178,805,831 (that’s sixty-two trillion four hundred ninety-one billion one hundred seventy-eight million eight hundred five thousand eight hundred thirty-one) isomers (somehow, it seems like more isomers if you spell the number out). The C(xy) compounds between C20 and C40 have numerous possible isomers as well and they increase logarithmically (see chart below) as the number of carbons increase. Not all of these hydrocarbons are in paraffin but these numbers should give you an idea of how chemically complicated a simple candle may be.
While this already seems like a brain-damaging subclause to our proceedings, the estimates for number of isomers for each number of carbon is actually more complicated than I am representing here. If you have further interest, you can take a look at this discussion. If not, let’s proceed.
There is a standard equation for calculating how much product results from combustion in oxygen of any hydrocarbon; it is:
where z = x + y/4.
This means that in cases where there are 20 carbons as for eicosane, the carbon dioxide and water molecules result in the following way:
2 C20H42(s) + 61 O2(g) → 40 CO2(g) + 42 H2O(g)
or… for each two molecules of n-eicosane (one of about 366 thousand isomers of eicosane) are consumed by combustion, sixty-one molecules of oxygen are consumed, thus producing 40 molecules of carbon dioxide and forty-two molecules of water.
The thing is that it is rare that anyone burns a candle or anything else in pure oxygen. When hydrocarbons are consumed in air, a messier equation obtains to the problem:
Note that carbon monoxide is produced, along with hydrogen gas and the more familiar carbon dioxide and water. This version of the equation is why it is critical to ensure adequate air supply when using a kerosene (or other hydrocarbon-based) space heater in a closed space; the amount of carbon monoxide goes up as the amount of oxygen available goes down. Carbon monoxide, a colorless and odorless gas, causes humans to fall asleep and die due to a special kind of asphyxiation caused by very strong binding of carbon monoxide to the iron atoms in your hemoglobin and myoglobin. Once that happens, those proteins cannot carry oxygen through your arteries and your body is “starved” of oxygen.
Okay, so hydrocarbons burn in air (n.b. there is also lots of nitrogen in air and that produces problematic by-products as well) and that means carbon monoxide, carbon dioxide, water, and hydrogen are produced, along with a substantial amount of particulate matter (e.g. particulate carbon and other solid carbon by-products), which ends up in our shared atmosphere (n.b. there is no “U.S.A. atmosphere” or “China atmosphere,” there is one planetary atmosphere). The most common liquid fuel currently consumed is octane but that is not consumed as pure octane, so there are other hydrocarbons and “stuff” consumed at the same time… in air… which produces problematic by-products.
Here’s a chart of how much world liquid fuel has been consumed and is projected for consumption PER DAY over the listed time period:
Yes, the chart does indicate that we consume between 94 and 96 million barrels of liquid fuel per day. One barrel of liquid fuel is equivalent to 0.1172 metric tons and a metric ton is 2,200 pounds (for the non-metricized readers). One barrel is 257.4 pounds of liquid fuel. If we are consuming (let’s be modest) 94 million barrels of liquid fuel per day (and let’s be factual) there are 365 days in a year, we are consuming 8,846,490,400,000 pounds of fuel per year. If we were to pretend that all of this were octane (which it isn’t) and all of that octane followed the simplest hydrocarbon-to-carbon dioxide equation provided above (which it doesn’t), we say that every two units of octane produces sixteen units of carbon dioxide. These don’t have the same mass, of course.
To make this simple, a gallon of gasoline weighs about 6 pounds. Each gallon of gasoline produces about 18 pounds of carbon dioxide (idealized as stated above). If we divide the number of pounds of liquid fuel consumed annually by 6, we will have an estimate of the number of pounds of carbon dioxide produced. Well, the number is:
(8,846,490,400,000 pounds of fuel per year)/(1 gallon/6 pounds) =
1,474,415,066,666.67 pounds of carbon dioxide/year
To do our numbers-into-language thing, that is one trillion four hundred seventy-four billion four hundred fifteen million sixty-six thousand six hundred sixty-seven (let’s round up, given the decimal figure) pounds of carbon dioxide produced from the aforementioned pounds of liquid fuel. Pretty incredible, right?
The bottom lines are these:
- we can’t breathe carbon dioxide (it chokes us)
- actual combustion produces lots of other by-products that are also not useful for human respiration and cause various respiratory illnesses (cancer, emphysema, asthma for starters)
- these numbers don’t include gaseous fuel like methane, ethane, propane, or butane (starting with pentane and going up to heptadecane (C17), the compounds are liquid at 25°C), which are also used as fuels.
- these numbers don’t include non-petroleum fuels such as ethanol, which is an oxygenated hydrocarbon but also produces all the by-products listed for hydrocarbons
- Our global economy is heavily dependent on consuming something that
- is finite in quantity and
- produces harmful by-products
- is going to go up in price as the amount available nears complete consumption
- We have not solved the equation for producing less carbon dioxide and less harmful by-products while maintaining our current lifestyles.
Okay, end of lesson. Talk amongst yourselves. This all needs to be solved.
Burn a candle while you’re at it. Couldn’t hurt (much).