Editor’s note: We are pleased to present a series adapted from biologist Michael Denton’s book, Fire-Maker: How Humans Were Designed to Harness Fire and Transform Our Planet, from Discovery Institute Press. Find the whole series here. Dr. Denton’s forthcoming book, The Miracle of the Cell, will be published in September. 

Given the great quantities of energy released in combustion — and given that our bodies are composed of reduced carbon compounds — a question arises, one that intrigued novelist Arthur C. Clarke: Why don’t we spontaneously combust, even at ambient temperatures, given the inherent thermodynamic energy of oxidations?¹ Indeed, why don’t forests do the same? 

A forest fire is ample proof of the enormous amounts of potential energy (thermodynamic) that “lie within.” It is claimed that in the great wet era — the Carboniferous, when the Earth was a massive swamp and the first amphibians crawled in the watery margins of the lakes and streams — oxygen levels reached 30 percent or even slightly higher. This is a huge proportion compared to today’s levels, and the evidence suggests that those “wet forests” burned in conflagrations of unimaginable ferocity.² The effects of these great conflagrations were attenuated only because the Earth was wet and most life was aquatic or lived on the margins of rivers and swamps. 

James Lovelock has pointed out that atmospheric levels of oxygen much above about 25 percent, let alone 30 percent,³ would cause raging conflagrations today even in tropical rain forests. So controlling fire in a normobaric atmosphere of more than 25 percent oxygen would likely be highly problematic. Current ambient levels close to 21 percent are just about ideal for controlled combustion: high enough to get a fire started, but not so high that the fire ignites spontaneously and spreads uncontrollably.⁴

Ideals Levels, but Why?

The reason why neither humans nor trees spontaneously combust at the current 21 percent oxygen levels (= pO₂ of 160 mm Hg) is because both the carbon atom and molecular oxygen (O₂, or dioxygen) are relatively inert at ambient temperatures because of their peculiar atomic structures, which greatly attenuates their reactivity.⁵

This attenuation of the reactivity of oxygen makes it possible to sustain the high metabolic rates of mammals on our planet. It means that atmospheric levels of 21 percent oxygen, which are required to supply air-breathing, energy-hungry organisms (like mammals and birds and flying insects, etc.) with sufficient oxygen to satisfy their metabolic needs,⁶ do not at the same time lead to spontaneous combustion. As Roman Boulatov comments: “The biosphere benefits greatly from this inertness of O₂ (dixoygen) as it allows the existence of highly reduced organic matter in an atmosphere rich in a powerful oxidant.”⁷

A Problem and Benefit for Life

Ironically, the chemical inertness of O₂ is a potential problem for life as well as a benefit. Boulatov continues, “[S]uch inertness also means that rapid aerobic oxidation will only occur if energy is put into the system to overcome the intrinsic kinetic barriers [e.g., heat is used to start a camp fire] or the reaction is catalyzed”⁸ by enzymes that contain either iron or copper ions within their active sites.⁹

 It is another element of fortuity in nature that the properties of the transition metal atoms, such as iron and copper, have just the right atomic characteristics to “gently” activate oxygen for chemical reactions. In fact, all the oxygen-handling enzymes in the body, even those not specifically involved in oxygen activation such as hemoglobin (which is involved in oxygen transport), make use of transition metal atoms. So the inertness of oxygen at ambient temperatures is rescued in the body by the unique properties of the transition metal atoms that activate it for energy generation in air-breathing organisms like ourselves, whose high metabolic rates and active lifestyles depend critically on the energy of oxidations.¹⁰ If not for our unique oxygen-handling capacities, we as carbon-based life forms dependent on oxidations for our metabolic energy would certainly not be here. 

Lucky for Us

In short, the inertness of dioxygen is clearly fit in several ways for air-breathing organisms obtaining their oxygen in gaseous form supplied from an atmosphere: It enables the energy of oxidations to be utilized in the body; it prevents us from spontaneously combusting; and it allows for the controlled utilization of fire.

It is worth noting that the inertness of oxygen at ambient temperatures is a fitness in nature particularly relevant for terrestrial, air-breathing organisms like ourselves, preventing spontaneous combustion and at the same time allowing for the mastery of fire. It does not apply to aquatic organisms that extract their supply of oxygen from water and are incapable of ever lighting a fire. And of course these characteristics are completely irrelevant to anaerobic bacteria and those extremophiles entombed in the crustal rocks, far removed from the concerns of life with oxygen.¹¹

Tomorrow, “Our Planet Is Just the Right Size.”

Notes

1. “This one mystery I’m asked about more than any other—spontaneous human combustion,” says Arthur C. Clarke, author of 2001, in the episode “The Burning Question” of his 1994 TV series Mysterious Universe. “…Yet some cases seem to defy explanation, and leave me with a creepy and very unscientific feeling. If there’s anything more to SHC, I simply don’t want to know.”
2. Nick Lane, Oxygen the Molecule That Made the World (Oxford; New York: Oxford University Press, 2002).
3. James Lovelock, Gaia: A New Look at Life on Earth (New York: Oxford University Press, 2000), Chapter Five. 
4. Ibid. 
5. Monika Green, H. Allen Hill, “The chemistry of dioxygen,” Methods Enzymology 105 (1984): 3−22
6. David C. Catling, Christopher R. Glein, Kevin J. Zahnle, Christopher P. McKay. “Why O2 Is Required by Complex Life on Habitable Planets and the Concept of Planetary ‘Oxygenation Time,’” Astrobiology 5, no. 3 (June 2005): 415–38. doi:10.1089/ast.2005.5.415.
7. Roman Boulatov, “Understanding the Reaction That Powers This World: Biomimetic Studies of Respiratory O2 Reduction by Cytochrome Oxidase,” Pure and Applied Chemistry 76, no. 2 (2004): 303–319, doi:10.1351/pac200476020303.
8. Ibid.
9. Corinna R. Hess, Richard W. D. Welford, and Judith P. Klinman, “Oxygen-Activating Enzymes, Chemistry of,” in Wiley Encyclopedia of Chemical Biology (Hoboken, NJ: John Wiley & Sons, Inc., 2008). http://doi.wiley.com/10.1002/9780470048672.wecb431. The authors comment: “Nature has developed a diverse array of catalysts to overcome this kinetic barrier. These dioxygen-activating enzymes are divided into two classes: oxygenases and oxidases. Oxygenases incorporate directly at least one atom from dioxygen into the organic products of their reaction. Oxidases couple the reduction of dioxygen with the oxidation of substrate. Typically, enzymes that react with dioxygen contain transition metal ions and/or conjugated organic molecules as cofactors.”
10. Catling, et al.
11. Thomas Gold, The Deep Hot Biosphere (New York: Copernicus, Springer-Verlag New York, Inc., 1999).