The Equations of Life

by Charles S. Cockell

  To answer this question and to explain why life has made the choices it has with all the elements it uses, we must know something about the structure of atoms from which life is assembled. By delving into the periodic table and the physics of atoms, we will find that extraordinary universal principles of physics ultimately lie at the heart of carbon as the favored element for constructing life forms.

  Developed in one of its first modern forms by Dmitri Mendeleev in 1869, the periodic table contains all the known elements, either known naturally or synthesized in the laboratory. At the core of every atom of every element is a nucleus, which contains protons, positively charged particles. Besides hydrogen (with only one proton), the nucleus of all other atoms has some neutrons too. These neutral particles play a part in binding the nucleus together. Elements are enumerated according to the number of protons they have, sometimes called the atomic number. So hydrogen, with one proton, is element number 1. It sits at the top left of the periodic table, and the somewhat awkwardly pronounced oganesson, element number 118, sits at the bottom right.

  Surrounding this small bundle of particles in the center of the atom are electrons, subatomic particles that also have something of a wavelike quality to them, like light. But unlike the protons, the electrons have a negative charge. Atoms are always neutral—they have no charge—so the positive charges of the protons must cancel out the negative charges of the electrons. In other words, the number of electrons in an atom must be the same as the number of protons.

  So far, we have a simple picture of elements increasing in atomic number from the top left to the bottom right of the periodic table. To create each element, one proton in the nucleus and one electron metaphorically in orbit around it are added sequentially, building the zoo of atoms from which the universe, and life, is made.

  There is a little problem in the view I have presented. We cannot just add electrons to the atom one by one in a growing crowd of the particles. Electrons hate being next to other electrons which are exactly the same, a little like identical twins dumped next to each other at a birthday party—twins who dislike being compared and who prefer their friends to treat them as distinct. Therefore, you can’t just stack electrons next to each other. This principle, that electrons, or all fermions, cannot occupy the same states, is called the Pauli exclusion principle, named for the inventor of this concept, Wolfgang Pauli, an Austrian-born physicist.

  What then, do we do with two electrons that are side by side in an atom and do not want to be identical? One property we can alter is their spin. If the electrons’ spin is in different directions (spin up and spin down), then they are now distinctive. Like the twins that have some distinguishing feature from which they can feel a semblance of individualism, the two electrons can now abide by the Pauli principle. However, this principle prevents us from adding a third electron, since there is no other property we can modify to make the third electron different. Like a metaphorical subatomic Noah’s Ark, the electrons are stacked into the atom two by two.

  As we add electrons to atoms, they occupy so-called orbitals, sometimes called shells. Each shell or orbital can contain two electrons or multiples of two electrons, ensuring that the Pauli exclusion principle is not violated.

  Once the stacking is completed, those outermost electrons that filed into the last orbitals are of singular importance because they are the part of the atom that will first come into contact with another atom; they will define the nature of any chemical bonds or whether atoms will react with one another at all. Atoms that have partly full electron orbitals like to gain or lose electrons to end up with a complete set of electron pairs; empty electron positions make atoms reactive.

  Pauli’s little rule explains why the noble gases, such as neon and argon, are famously inert. They have their outermost electron shells full to the brim with electron pairs—four pairs of two electrons, with no room to spare—meaning there are no spaces left to accept electrons from other atoms or to participate in exciting chemical reactions. The noble gases are left conservatively unreactive.

  This stacking of electrons in atoms explains how the elements, from 1 to 118, are arranged in the periodic table. Each column of the table contains the atoms that have the same number of electrons in their outermost shell. This means that within each column, because they have the same number of outer electrons, the elements have very similar chemical properties. So now you can see that the properties of atoms and the way they shape the material world around us are decided by the way the electrons are all stacked. That is determined by a simple physical principle: the Pauli exclusion principle.

  Let us return to life and consider the element at the core of most of its molecules: carbon. It has six electrons. Those six electrons must be stacked in a way to keep Pauli satisfied. Two electrons sit in the so-called 1s orbital, the lowest orbital. Two electrons can then be stacked into the next orbital up (the 2s orbital). The remaining two are in another orbital at the same level, the 2p orbital.

  What about the Horta? The speculative creatures are made of silicon, which is in the same group of the periodic table as carbon is, but one row down. Silicon contains fourteen electrons. How are they stacked? Two electrons are in the 1s orbital, and two in the 2s orbital, like carbon. Six are then stacked in the three 2p suborbitals. Two are placed in the next orbital up, the 3s orbital, and then finally two in the 3p orbital. Although silicon has more electrons than carbon does, the two elements’ outermost shells are very similar—two electrons in an s orbital and two in a p. This similarity explains why carbon and silicon have similar chemical properties and why the Horta came to exist in our minds.

  Now that we have a grasp of one fundamental principle that lies at the heart of biology—at its lowest level of its hierarchy, the atomic level and its constituent subatomic components—let us explore a little more what makes carbon a good building block for life and whether silicon might suffice, too.

  Carbon is just the right size. In its outermost shells, the electrons that stand ready to pair up with electrons in other atoms and form bonds, thereby forming molecules, are bound close enough to the nucleus that they hold on tight, meaning the links are strong. They are not so far away that they are pulled off the atom easily, which would make the bonds liable to break. Life must be able to build molecules that, like DNA, will be stable, but it must also be able to pull apart old molecules to make new ones without using vast amounts of energy. Carbon fits the bill.

  The electrons in the outermost orbitals, those electrons in 2p and the other pair in 2s, love to pair up with electrons from other atoms and form bonds. In an especially common reaction for a carbon atom, one of its electrons connects with the single electron in hydrogen to form a carbon-hydrogen bond, a union that decorates all manner of life’s molecules. Carbon can form bonds with other carbon atoms, and with sulfur, phosphorus, oxygen, and nitrogen too. These bonds have similar strengths, so carbon needs little energy to switch between these different atoms. The atom has other arrangements. It can form double bonds. The two electrons in the 2p orbital can pair up with two electrons in the 2p orbital of another carbon atom and form the double linkage. This capacity, along with an ability to form triple bonds, adds to the welter of carbon-containing molecules.

  Resulting from all this versatility and enthusiasm to form bonds is an impressive diversity of chains, rings, and other structures: from the simple gas methane, made of just one carbon atom bound to four hydrogen atoms, to the amazingly long molecule DNA, which unraveled is a full two meters long in a human! It is therefore natural for anyone to ask, when confronted by this flexibility in the assembly of a range of molecules, if other elements could do the same. Silicon is an obvious contender, and as the second-most abundant element on Earth after oxygen, it might well look like a pretty good candidate.

  Despite the similar electron configuration on the surface, there is one crucial difference between silicon and carbon. As noted, silicon has fourteen electrons stacked up, in contrast to carbon’
s six, which means that silicon’s outer electrons are further away from the nucleus and less tightly bound than carbon’s outer electrons. A consequence of silicon’s more lightly bound electrons is that its bonds with other molecules tend to be weaker than carbon’s. The silicon-silicon bond is about half as strong as the carbon-carbon bond, meaning that rarely in nature can you find more than three silicon atoms bound side by side. As a result, there is little chance that silicon can build all those complex chains and rings that we find in carbon-based life, where many dozens of carbon atoms can be joined in chains. The electrons, being less tightly bound to the nucleus, are more apt to be snatched up by other atoms or themselves to pair up with other electrons, making the atom more reactive. Some bonds that silicon forms are very unstable. Silane (SiH4), an analogous molecule to the biologically important gas methane (CH4), spontaneously combusts at room temperature.

  However, silicon has another Achilles’ heel. A carbon atom, when it binds to an oxygen atom, can form a double bond. With two oxygen atoms, the result is the very versatile gas carbon dioxide, which is the raw material of photosynthesis. However, silicon, because of its larger size, cannot so easily form a double bond with oxygen and instead forms four single bonds that are more comfortably distributed around the larger atom. These oxygen atoms still have a single bond to spare, and they use this to bind to another silicon atom. The result? A giant network of silicon and oxygen bonds linked to form a great grid. And this grid is very familiar to you and me. It is the structure of the silicates, the material that makes glasses, minerals, and rocks. Unfortunately, unlike many other silicon compounds, they are so stable that once silicon is locked up in these structures, it obstinately stays there. Rocks are one of the most visible reasons why silicon-based life is unlikely.

  These rocky silicates can be found in a dazzling variety. In a crude way, they recapitulate the enormous treasure chest of carbon compounds. But silicates are the stuff of rocks, not biochemistry. Their networks make them inert, so unreactive that silicate ceramics are used as heat shields to protect spacecraft as they enter the Earth’s atmosphere; the intense temperatures, rising well above a thousand degrees Celsius, are still unable to cajole the structure of the material to do something interesting.

  Although most of the silicon on our planet may be locked up in generally unreactive silicates, life is by no means devoid of the element. Diatoms, algae that inhabit the oceans and freshwater rivers, lakes, and ponds, protect themselves inside a frustule, an ornate shell made of silica (silicon dioxide). These photosynthesizing microbes achieve a beautiful diversity in their forms, including stars, barrels, and boatlike shapes. Plants also gather up and use silica. In some, the amount of the substance may be up to a tenth of the plant’s total mass! Silicon is readily absorbed from soil as silicic acid and is thought to play a part in growth, mechanical strength, and resistance to fungal diseases. It is prominent in phytoliths, silica structures that are formed in cells and that aid in the plant’s rigidity, necessary for upward growth against gravity. Silica structures, known as spicules, are even found as a primitive skeleton in certain marine sponges, organisms that belong to some of the earliest multicellular organisms on Earth.

  No sensible scientist would discount silicon as a basis for life. Even on Earth, where silicates form 90 percent of the crust, the element need not be entirely bound with oxygen in silicates. The silicon and carbon compound silicon carbide (SiC) occurs naturally. In the interstellar medium, many silicon compounds such as SiN (silicon mononitride), SiCN (silicon cyanide), and SiS (silicon monosulfide) have been observed, showing that on the universal scale, silicon can make some unusual compounds. We have a certain bias since we know so much more about carbon chemistry than we do about silicon chemistry. As we delve further into silicon chemistry, we come across some surprises. The atom seems to form a colorful variety of compounds in concert with carbon—the organosilicon compounds, some of which form chained arrangements. Perhaps a black-and-white view of the two elements misses the potential for some sort of hybrid carbon-silicon-based life form.

  Give it a chance, and silicon can form more-promising products. Among its family of structures are the cage-like molecules with the tongue-twisting name of silsesquioxanes. All sorts of structures can be added to these molecules’ core to produce a magnificent collection of other molecules. Other silicon compounds, under just the right laboratory conditions, can be enticed to form silicon chains with over twenty consecutive atoms, like the long-chained compounds that make up the molecules of living things.

  Although these amazing jaunts through silicon chemistry show that there are complex silicon compounds that have astonishing multeity, life has not been idle. It has evolved to test out this element in many functions, but it has not yet, as far as we know, used silicon to build the major molecules of life to such a pervasive extent that we would describe an organism as silicon based. Silicon-filled plants still have cells constructed with sugars, proteins, and lipids, the stuff of carbon-based chemistry. Tellingly, when life gets hold of silicon, cells do rocklike things with it—building siliceous structural support materials such as phytoliths and spicules. While perhaps life’s structural use of silicon is a vestige of Earth’s evolutionary history of life, which chose carbon, if organisms found some benefit to using silicon in many compounds that enhanced their chances of survival, they would use it. Earth’s evolutionary experiment shows that under the conditions associated with our planet, carbon trumps silicon in almost all biochemistry.

  The other elements in group 14, the same group that carbon and silicon belong to, suffer problems as the atoms grow in size. Germanium is the next element down, but germanium life forms have never been entertained. As far as we know, this element is also unable to produce the range of chemical compounds useful for building living systems. Onward down into group 14, tin or lead Horta seem to have even less chemical evidence to support their existence.

  For all the pushing and pulling we can do with the periodic table to find elements that might look like good choices for fashioning life, carbon remains far and away the element with the largest and most diverse number of molecules in its repertoire of bonding possibilities. It is likely that the processes of life elsewhere in the universe would converge on this element as the basic elemental building block of life. And as described earlier, carbon is the best choice because of the Pauli exclusion principle, a universal principle operating at the quantum level and laying down rules for how electrons are stacked within atoms.

  Those people with a healthy skepticism might still be unconvinced. Could other life forms modify not only the core atoms that they use, but also the solvent in which life operates? Maybe our assumption that carbon chemistry is linked to water limits our capacity to imagine alternatives. Should we consider a different union between the liquid and the central elements of life? An imaginative, if highly unfamiliar suggestion is that silicon-based life could originate and evolve in liquid nitrogen. The liquid nitrogen would offer sufficiently cold temperatures for complex and otherwise generally unstable silicon compounds, such as silanes and silanol compounds, the latter analogous to alcohols on our own world, to remain stable.

  A wild and extraordinary geological cycle can be concocted from these suggestions. Reactions of silica in the rocks of a planetary core with carbon dioxide, ammonia, and other compounds would produce silanes and silanols. These would eventually be transported into the liquid nitrogen ocean, where they could participate in further chemistry, providing the basis of a silicon-based life. A location for such a novel type of biology has been suggested to be Neptune’s moon Triton, an ice-covered world with nitrogen geysers on its surface, possibly erupting from deep, buried, frigid liquid nitrogen just beneath its surface. However, any place with some rocks and liquid nitrogen will do for this bizarre system of life.

  These types of exotic chemistry and solvent combinations are even more difficult to assess than a relatively simple swapping out of one atom for another, since the
y stray into the realms of chemistry we know little about. The full capacities of silicon chemistry in liquid nitrogen are quite unknown, and we cannot rule out life in such circumstances in light of our knowledge of chemistry alone.

  There may be good reasons, though, for seeing carbon in a positive universal light, even with these interesting alternatives on our minds. Not only does carbon have promising atomic physics for building complex life forms, but this propensity to form multitudinous molecules ensures that carbon molecules are abundant in the universe, making it likely that other examples of evolution, if they exist, would from their earliest stages find carbon molecules the most readily available larder of complexity.

  Look up into the sky on any clear night, and you see a universe that billions of pairs of eyes have gazed upon throughout the history of our civilization. A canvas of black interspersed by the glinting white points of celestial bodies. Every now and then, their unchanging positions are interrupted by a comet, the bright glow of a supernova, or the nightly streaks of debris burning through our atmosphere as shooting stars, but otherwise the night sky seems immutable in human life spans.

  This view of space as an endless landscape of emptiness is not inaccurate, at least compared with the rich variety of matter that packs onto our small world. However, to believe that the vastness of the universe is barren, a view that has dominated our collective consciousness since we first realized those points of light were stars and the blackness between them the rest of the vacuum of the cosmos, would be to overlook the breathtaking complexity of chemistry that goes on in this apparent void.

  In the beginning, when the big bang heralded the beginning of the universe, things were simpler. As temperatures dropped, chemistry was confined to a few reactions between hydrogen, helium, lithium, and their ions, with some electrons and radiation thrown in for good measure, a basic playground of elemental rearrangements. Then the first swirls of gas gravitationally collapsed to a sufficient density to trigger the fusion reactions of stars. Within these glowing balls, the joining of hydrogen atoms into heavier elements, carbon included, could occur.