The Equations of Life

by Charles S. Cockell

  The Miller-Urey experiment established that organic compounds can be formed on the surface of a young planet, not just in space. So here we have it—organic molecules forming in interstellar space and on the surface of a planet. From above and below, young planets are crucibles in which a selection of the enormous variety of organic compounds that exist can collect and become available for life.

  The tendency for complicated organic chemistry to happen in places with some energy and basic ingredients is spectacularly demonstrated in all the organic chemistry occurring on Titan. Its methane lakes are the source of a cosmic factory of organic molecules. Its brown atmospheric haze is a product of atmospheric methane reacting with UV radiation high in its atmosphere, breaking down to form radicals that subsequently reform into ethane and complex chains of organic compounds, some of which drift in the upper atmosphere, producing the unusual hue.

  Much of this material rains down on Titan’s surface to produce vast deserts of organic compounds. In some regions of the moon, dunes made of complex carbon compounds ripple for hundreds of kilometers across the surface and up to a hundred meters high. If what is happening on the surface of this moon has been going on since the beginning of the Solar System, there would theoretically be a layer of ethane, the carbon molecule with the structure C2H6 and the precursor to many more-complex interesting molecules, about six hundred meters thick!

  Are there building blocks of life mixed in with all this organic material on Titan? We do not currently know. Future robotic missions to Titan may answer this question, but whatever the answer, it does not impinge on the simple observation that on a moon with some methane and some radiation, vast dunes, lakes, and atmospheric hazes of organic compounds are being fabricated. Complex carbon chemistry is a planetary rash.

  Some people think it uncanny that on Earth, conditions are just right for carbon chemistry to produce life. What we have seen suggests exactly the opposite. Everywhere, under conditions vastly different from those on Earth, carbon forms a huge array of versatile, reactive compounds. We find that, under a great range of temperature, pressure, and radiation conditions, of all the elements in the periodic table, carbon produces the greatest diversity of molecules. Yes, certainly, we find other compounds; even silicon-carbon bonds and silicon-nitrogen bonds detected in interstellar space suggest that extraterrestrial conditions can produce strange and intriguing new unions between elements rare on Earth. They give us pause for thought about the limitations of our chemical knowledge. However, glaring indefatigably through all these data is the array of carbon-based compounds that suggest that the variety of such structures on Earth is not unusual. What Earth may well have provided is the environment for these compounds to form long chains that then evolved into self-replicating entities. That step may require particular conditions not easily achieved in gas clouds and frozen dunes. But that event occurred in a universe where a carbon-based chemistry is the norm everywhere.

  A more fascinating story develops when we expand our horizons beyond carbon to think about the other elements known to be spread through life. To build a life form, we need more than just carbon atoms. Of all the atoms available, five are ubiquitously attached to carbon to make more-complex arrangements: hydrogen, nitrogen, oxygen, phosphorus, and sulfur. These chemicals, along with carbon, are sometimes remembered with an unmemorable mnemonic, CHNOPS. Why these atoms? Is their ubiquity in life also a product of simple physical considerations?

  These elements, too, just like carbon, owe their behavior to the Pauli exclusion principle. Hydrogen is always ready to bond with other atoms from which it can gain just a single electron to get a complete pair and fill its only electron shell. Throughout life, we find it bound to carbon. Hydrogen can be crudely regarded as a sort of mopper-up of single electrons that are available. That explains its appearance in everything biological.

  The other four elements that, like nuts and bolts, hold living things together through their carbon networks—nitrogen, oxygen, phosphorus, and sulfur—intriguingly occupy a small quadrangle of the periodic table, huddled together near carbon. At the grand chemical scale, we can fathom this unusual affinity. All four have incomplete electron orbitals and will take part in bonding with other atoms to fill them. They inhabit a space in the periodic table where the size of the atoms is just right. Because the electrons can form bonds that do not need too much energy to break up, these elements are useful in the constant assembly and disassembly that characterizes the building and growth of living things. Nevertheless, the electrons in these atoms are attached tightly enough to make their links generally stable. In broad terms, these four elements are just very good at forming a diversity of compounds with multifarious properties. When they form bonds with carbon, they contribute to the imposing variety of chemical reactions required to be a successful reproducing and evolving life form.

  On the face of it, nitrogen does not look like a very promising element for life. It makes an incredibly strong triple bond with another nitrogen atom to form nitrogen gas, N2, which makes up a prodigious 78.1 percent of the Earth’s atmosphere. However, once released from this chemical prison by catalysts in so-called nitrogen-fixing microbes or by nonbiological processes like lightning, the nitrogen is available to form a variety of useful bonds with carbon. One of its stable configurations is to sit between two carbon atoms and form the linchpin of the peptide bond, the linkage that holds individual amino acids together to form proteins. All amino acids contain nitrogen, which allows them to string together in this way. Nitrogen also turns out to be good at sitting between carbon atoms to make rings, and so it can be found in many of the prominent ring-shaped molecules in life, including the DNA base pairs. The nitrogen in these nucleic acids is the hookup to the sugars in the backbone, helping to hold the whole edifice of the genetic code together.

  Let’s now take a small step to the right in the periodic table and think about oxygen. This ubiquitous gas, so essential for animal life, sits neatly up against nitrogen in the table. Oxygen atoms perform a somewhat similar role to nitrogen, connecting carbon atoms in rings and forming attachments between carbon-containing molecules such as sugars, allowing long chains of sugar molecules, or carbohydrates, to form. Oxygen-containing sugars form part of the backbone of the crucial nucleic acids. We find oxygen in an array of organic molecules such as carboxylic acids, which participate in the synthesis of more-complex molecules such as proteins.

  Phosphorus and sulfur, the two other molecules of our quad, sit below nitrogen and oxygen, respectively, and so they are larger with their extra electrons.

  Familiar to most of us as the combustible material at the end of matches, phosphorus has insinuated itself into many of life’s key molecules. Because it is larger than the other CHNOPS elements and its outer electrons form and break bonds more easily, phosphorus turns out to be a key component of many of life’s energy-requiring reactions. Attach oxygen to it, and now you have a bond that can be broken to rapidly release energy in hydrolysis reactions where it is needed. The molecule ATP, with three phosphorus atoms strung together between oxygen atoms, has become one of the quintessential energy-storage molecules in all life on Earth, like a miniature battery for life.

  The enormous versatility of phosphorus extends to the very structure of cells themselves. The phosphorus atom is found at the end of lipids, the long-chained carbon compounds that come together to make cell membranes. Even in the genetic code, this element is liberally distributed. Strung down the backbone of DNA, the phosphorus atoms form the links between the sugars, holding the structure together and helping provide longevity. The negatively charged oxygen atoms, found dangling from the phosphorus, give DNA a negative charge, which stops the molecule from leaving the cell since it is repelled by the negative charges on the inside of the lipid membrane. These negative charges have a second role: they help prevent DNA from being hydrolyzed, making the molecule much more stable.

  To the right of phosphorus is sulfur, the biblical brimstone, or “bur
ning stone,” the yellow substance that adorns and drapes the vents and cauldrons of active volcanoes. Like phosphorus, its elemental association with fire and violence belies its beneficent use in life. The element is found in proteins. Two sulfur-containing amino acids in different parts of the long chain of amino acids that make up a protein can come together to form a disulfide bridge—two sulfur atoms linked side by side. These bridges help bolt together the three-dimensional structure of a protein as they link different parts of the chains into a well-defined shape that can carry out catalytic reactions in the cell.

  These are just a few of the many uses of the four of the six CHNOPS elements necessary for life, but they illustrate the adaptability of these atoms and some of their different characteristics that find uses in cells as part of the machinery of life. Their deployment in many complex long-chained molecules shows just how useful they are.

  Nitrogen, oxygen, phosphorus, and sulfur seem to have some convincing and useful applications in life. Evolution has used them often, accounting for their ubiquity and their elevated status to CHNOPS elements, but what about the alternative elements? Can we rule them out?

  To the right of our chemically familial set of quads, and sitting next to oxygen, is fluorine. The element gained its fame in the process of fluoridation, a postwar attempt to reduce tooth decay by adding small quantities of fluoride to public water supplies. However, aside from this largely quiet entrance onto the stage of human society, fluorine is generally not used extensively by life. Its electron shell is almost full—with seven electrons in its outer electron layer, it would like eight electrons to make four pairs. Those seven are tightly bound to the nucleus, and the fluorine atom is highly desirous of that last electron in the set, like a small child who runs around in gleeful acquisitiveness to get the last color for a complete set of colored toys. This property makes it reactive, and when it binds to other atoms, it will not let go easily. The carbon-fluorine bond is the second-strongest bond in organic chemistry. Carbon-fluorine bonds are just too inert and strong to find much use in life.

  But the fluorine atom is not a biological pariah. In the tropics, numerous plants and microbes use fluorine compounds as poisons to deter predators. As is the case for all the other non-CHNOPS elements throughout the periodic table, if some specific chemical property of a compound containing an element can find use in the struggle for survival, life will evolve to use it. The point is that the chemical properties of fluorine, accounted for by its electron structure, make it of limited use. It cannot be a universally used atom in life.

  Underneath fluorine, the element chlorine suffers from similar problems. Although its greater number of electrons makes the atom larger and less desirous of the last electron, it still likes to get that final electron in the set, a propensity that gives the element sufficient reactivity to earn its place in the chlorine-based bleaches deployed in your bathroom to kill unwanted microbes. Not that it is discarded by life, either. It can be found in cells and carries out functions such as balancing the concentrations of different ions, but its chemical properties limit the scope of its use.

  And beneath the nitrogen, oxygen, phosphorus, and sulfur? Below our quad are two intriguing elements, arsenic below phosphorus and selenium underneath sulfur. Both are found widely in living things and show us again that life never treats an element as an outcast. That said, the larger size of the arsenic and selenium atoms compared with their smaller CHNOPS counterparts makes their electrons less well bound; bonds with other atoms fall apart more easily. Nevertheless, this loose-binding property has not dampened enthusiasm for their possible role in alternative life forms.

  Switch out phosphorus with arsenic below it on the periodic table, and you have invoked the weird and alien-like possibility of microbes that contain arsenic in their major molecules. In an article published in the journal Science in 2010, the astonishing claim was made that a bacterium isolated from the alkaline and arsenic-containing Mono Lake in California had replaced the phosphorus in its DNA with arsenic. Much ado resulted, launched with a press conference and wild enthusiasm for this new turn in the biochemistry of life. However, in a little less than a few days, the incredulity of the scientific community was being expressed online and in the press. Why? The tendency of arsenic-containing compounds to rapidly hydrolyze, or fall apart in water, because of that bigger atom, made it unlikely that the DNA molecule would hold together. The microbe, after further investigation, turned out to be a bacterium that uses phosphate in its DNA, just another member of the family tree of life that we know and love.

  Some people might be tempted to say that this is how science works. Data are collected, claims are made, and sometimes they are refuted. This is how the scientific method advances the state of knowledge, however unpleasant the process can sometimes feel. However, even when this paper was published, the doubts about its credibility accounted for the rapid counterclaims that emerged. The estimated half-life of DNA bonds containing arsenate ions is about 0.06 seconds. If you replace arsenate with phosphate, the half-life jumps to about 30 million years. It would take very special circumstances or a lot of energy for a microbe to hold together a biochemistry with DNA in which the phosphorus atoms were replaced with arsenic, and this is true of many other phosphorus-containing molecules. The arsenic-using bacterium seemed unlikely from the start.

  To put a positive spin on this episode, it shows that the elements that flank the CHNOPS elements have such similar chemical properties that it seems plausible that they could, under particular conditions, replace each other—so plausible that scientists could be enticed to think that a microbe could replace the phosphorus in its DNA with arsenic. Like all research, it stimulated more studies of arsenic in life. Out of any scientific claim, there is always the possibility of advances.

  Nevertheless, despite this setback, we do know that life can contain arsenic. Its uses are enigmatic, but sugars containing arsenic have been found in some seaweeds, and arsenobetaine, an arsenic-containing molecule, is found in some fish, some algae, and even certain lobsters. Generally, though, arsenic is toxic to life. Its great propensity to share electrons means that it interferes with other molecules, reacting with them and disrupting metabolism. Many life forms have pathways to reduce or remove the toxic effects of the element.

  Arsenic’s next-door neighbor to its right, selenium, could plausibly replace sulfur, the element above it. This prediction is borne out in life. The element is found in one of life’s stranger amino acids. The so-called twenty-first amino acid of life, selenocysteine, is diligently included in some proteins. The energetic cost and modification to the genetic code required to include this amino acid means that its incorporation into living things is not a mere fluke of an exchange with sulfur. Selenium must perform a vital task. Some proteins in which it is found, such as glutathione reductases, prevent damage caused by oxygen radicals, which are reactive and potentially biologically damaging states of the oxygen atom. As an atom larger than sulfur, selenium can hand over its electrons more easily. This property plays a role in its ability to neutralize, if you will, the damaging free electrons in oxygen radicals. Once the selenium atoms have carried out this important function, they are more easily returned to their original state to carry out similar reactions. This reversibility, again because selenium can gain and lose electrons more easily than sulfur can, makes selenium useful in these roles. Moreover, with selenium, proteins seem to develop more resistance to oxidation, the process of losing electrons, caused by various types of chemical attack.

  Again, we see the same pattern: arsenic and selenium are not entirely rejected by life, but their size and the behavior of their electrons put them just outside that ability to be useful in many contexts. They have a specialist use in a few situations.

  To complete this tour around the fringes of the CHNOPS elements, let us examine one element near carbon in the periodic table—an element that we have so far neglected. Boron, an interesting little atom with three electrons in its ou
ter orbitals, can share those electrons with other elements. It can form bonds with nitrogen and, in so doing, generates compounds like borazine, which is analogous to the ringed carbon compound benzene. Boron lacks the chemical versatility of carbon, but like the other elements that surround the CHNOPS elements in the periodic table, it finds biological uses. It is an essential trace element in many plants, microbes, and animals and is thought to stabilize cell membranes and transport sugars. This is no ephemeral role: boron deficiency is one of the major trace-element deficiencies in agriculture, causing crop failures in plants from apples to cabbages. Our knowledge of the diverse roles of boron in biology is still in its infancy.