CHAPTER 9
WATER, THE LIQUID OF LIFE
SAMUEL TAYLOR COLERIDGE WAS not an astrobiologist, but his mariner’s observation that there was “water, water, everywhere” was about as important an observation as you could make about this most fundamental requirement for life.
There is a lot of water on Earth, about 1.4 billion cubic kilometers, or in more mundane terms but even grander numbers, about 560 trillion Olympic-sized swimming pools. Only about 0.007 percent of this water, freshwater, is actually used by you and me. The rest of it is in seawater, estuaries, marshes, swamps, and deep underground, inaccessible to humans but available to much of the rest of the biosphere, such as the microbes.
Carrying out the chemical reactions for life in a liquid makes sense. In a fluid, molecules can be brought together close enough to carry out reactions. Importantly, many millions of molecules can move around and meet in many combinations, aligning to do chemistry and to drive the complexity of pathways in living things. Such interaction is usually difficult to achieve in a diffuse gas cloud or in a solid. In a solid, molecules and atoms are generally rigid and cannot move around easily. In gases, they are often too far apart, in other words, too diffuse.
We might argue, with a little imagination, that in a gas, molecules would just react slowly and meet each other infrequently. The quixotic intelligent interstellar cloud, the product of imaginative science fiction, for instance, is perhaps just a bit cumbersome and not very chatty. However, the diffuse nature of the molecules and atoms in such clouds makes it unlikely that such a form could forge a self-replicating system that would evolve or be sustained over long time spans, let alone over the lifetime of a galaxy.
A compelling question—must life use water as a solvent?—has intrigued biologists for decades. In thinking about the answer, we consider contingency in the most basic requirements for living things—the liquid in which its parts are assembled. We continue our odyssey into the physical principles that shape the molecular level of life. Although water may look simple, a mere atom of oxygen bolted on to two atoms of hydrogen, this image belies its essential role in living things and the stunning variety of physics that explains life’s attachment to the substance.
We know of no single organism that can be active without water, and we know of no form of life that can use an alternative solvent to do the bulk of its essential chemical reactions. The question is whether this requirement for water results from one very specific set of evolutionary conditions or whether it derives from something more fundamental.
It has long been recognized that water has some very unusual properties. One of the most notable to you and me is that when it freezes, ice floats on water, since the frozen water becomes less dense, a simple observation that you can verify by observing ice cubes in your chilled drink. This property is strange, but not unique; silicon displays a similar behavior at a pressure of about twenty gigapascals. Most liquids, however, when they become solid, become denser and sink in their corresponding liquids. Water has this apparently anomalous behavior because the individual molecules in the liquid link up with each other in hydrogen bonds. The oxygen atoms in one water molecule line up with hydrogen atoms in other molecules, a result of the polar, or bar-magnet-like, qualities of water. In the liquid state, water molecules are agile and move around freely. They can get close to each other and twist and turn to fit in the nooks and crannies. However, when frozen, those hydrogen bonds rigidify, forming a well-ordered network that, because of its regular structure, takes up more space than the liquid does. With its more spaced-out structure, ice becomes less dense than liquid water and it floats.
Because of this unusual behavior, ice remains on the surface of a frozen winter pond, leaving the fish underneath protected inside their watery habitat while everything above them is solid. The icy roof traps the heat beneath, slowing the further freezing of the pond and allowing the fish to enjoy a bird-free existence, at least until spring. These village pond observations have led many to gasp in amazement that the physics of water seems so well tuned to life, for if ice sank, the village pond would freeze from the bottom up, killing the fish. However, we should not jump to conclusions on the vital nature of this uncanny and apparently life-supporting property of water.
In the forests of North America lives an enchanting animal, the wood frog, Lithobates sylvaticus. It inhabits the undergrowth and, to the casual observer, looks like nothing special. However, come winter, the little creature has a fiendish trick. When the winter frost arrives, the frog buries itself underground in leaf litter and soil and, in a feat of biochemical wizardry, produces the sugar glucose in its bloodstream. The sugar prevents the blood from freezing, curtailing the formation of ice crystals, whose long, slender shapes might otherwise rupture the blood vessels and damage the frog. When spring returns, the frog warms up and hops off into the undergrowth, unfazed by the turn of events.
The ingenious wood frog is an example of how we should be cautious in how we view the world. The fish in the frozen village pond might well suggest that the unusual properties of water are just right for life, but the wood frog shows us that if life evolved in a fluid that froze solid during the winter, life could probably adapt to these conditions. The properties of water are not adapted to life; life adapts to its surrounding chemical and physical conditions, including the fluid in which life happens to exist. However, this observation still does not answer the question of whether water has properties that make it a unique solvent to provide the crucible for living things to emerge.
Some aspects of water make it far from ideal for life. If we dig hard enough, we can even find properties that are deleterious. The substance looks innocuous enough in your glass, but water is not inert and it has an unpleasant ability to react with some of the key molecules of life. Hydrolysis reactions, as you may have gathered from the root of this word (from the Latin hydro-, or water), are reactions in which water can cause chemical changes.
In the liquid state, water is not merely the familiar formula H2O, but breaks up to form hydroxide ions, OH−, and hydronium ions, H3O+, which are protons (H+) bound to water:
2H2O ↔ H3O+ + OH−
The ions formed from the dissociation of water in this way can attack the long chains of life. From nucleic acids to sugars, hydrolysis reactions can cause these essential molecules to break apart, forcing life to use energy to constantly repair and rebuild itself against this damage.
Water may not be perfection, and with this contrarian outlook, we can certainly find aspects that speak against it. However, these minor quibbles aside, it has remarkable properties used in life.
Because its constituent atoms are slightly charged, or polar, liquid water can dissolve a wide range of small and large molecules—important for dissolving all those substances involved in the complex cascade of metabolic processes of life, from ions to amino acids.
Proteins—the diverse set of molecules that include the biological catalysts, enzymes, and many other pieces of biochemical machinery—have an uncanny and astonishing diversity of uses. Here, we see the true character of water and can view, in its full glory, what makes water such a good playground for the chemistry of life.
By binding to the outside of proteins, water molecules help keep them flexible, enabling them to move around sufficiently to take up the ingredients of the chemical reactions they drive as biological catalysts, but ensuring that they have enough rigidity to fold properly, maintaining their integrity. Strangely, in this role, water, often thought of as essential to maintaining stability, actually assists in destabilizing the protein just enough to encourage fluidity, showing the fine balance it oversees in life.
In other proteins, water molecules shield amino acids, preventing them from binding too strongly to other amino acids. This behavior, although apparently preventing stabilizing bonds from forming, again encourages just enough instability for the protein to remain flexible.
Even stranger collusions between water and proteins have be
en reported. The water molecules attached to the surface of a protein, because of their hydrogen-bonded network, become a “shell” of tightly bound molecules encapsulating the molecule. It is a physical state a little like glass. This behavior too plays a vital role in holding proteins together, but also in ensuring the easy movement of many of them.
In all these astonishing ways, water helps proteins to fold and enables their floppy chains of amino acids to coalesce in the right way. But it even goes further than this. Water can become part of the very structure of proteins, defining the shape and function of the overall molecule. By binding to the interior of the so-called active site, the region where the chemical reactions occur, water can link up with the incoming molecules, facilitating the catalytic roles that proteins perform. Water molecules are very much part of the way that many proteins work.
Not content with just getting into proteins, water deftly instantiates itself into the very code of life. How water molecules bind to DNA depends on the sequence of nucleotide letters within the DNA itself, so if the water molecules bound to the DNA then meet up with other parts of the DNA or other molecules in the cell, they are thought to mediate biochemical alterations related to the DNA code beneath them. This arrangement allows the genetic code to be read in an entirely unconventional way through the medium of water.
Water’s roles in cell biology go beyond helping structures form and orchestrating important reactions; cells also take advantage of the liquid’s ability to move electrons and protons around. Long chains of water, hydrogen bonded and behaving like wires, conduct protons around in bacteriorhodopsin, the molecule responsible for photosynthesis in some bacteria, allowing them to gather energy. In this clever arrangement, we see how the movement of subatomic particles through water is vital for some living things to acquire energy.
It is easy to dismiss some of this evidence as fascinating but a distraction from other possibilities. Much is made of the fact that some proteins can operate in non-water liquids such as the organic solvent benzene; this capability of proteins suggests that biochemistry could evolve in other fluids. However, most of these proteins must first be folded in water before they can do tricks in non-water solvents. The ability of some proteins to operate in organic solvents does not demonstrate that a whole biochemistry, with its many molecular interactions, could occur in other liquids or that even if it could, it is evolutionarily likely to happen elsewhere. Even when proteins are operating in nonaqueous alternatives, water is often still bound to these proteins and involved in their structural arrangement.
These different uses of water, in their intricate and manifold varieties, have driven home the point that we cannot simply think of life wallowing around in its solvent, but that the fluid is a fundamental part of the biochemistry of life. Life and its liquid are interwoven with such complexity and subtlety in so many ways that water is part of the machinery, not merely a medium in which other life-giving reactions happen to occur.
The impressive versatility of water, its diverse personalities that span across electron shuttle to proton wire, from hydrogen-bonded network to purveyor of molecular rigidity and flexibility, suggests that it may be unique in its capacity to integrate into, and play a major role in, a self-replicating, evolving living system.
Despite the growing list of water’s impressive attributes, what we know about other liquids should give us pause for thought. One most popular solvent thought to be a possible alternative for water in life is ammonia (NH3). At the equivalent of one Earth atmospheric pressure, its liquid temperature range is −78 to −33°C, but if you pressurize it, the boiling point can be raised to about 100°C, like the wide temperature range for water. Ammonia, like water, can dissolve many small molecules and ionic compounds too. The liquid might offer the potential of an environment for life where cold liquid ammonia solutions are thought to exist, such as in the deep subsurface of Saturn’s moon Titan, in the atmospheres of gas giants such as Jupiter, or maybe in the oceans of icy moons. However, there the similarities with water end.
An essential characteristic of life is the ability to compartmentalize molecules from the outside environment using membranes. Liquid ammonia does not support the spontaneous formation of membranes in the same way that water can, although at low temperatures, hydrocarbons, including some lipids, can be separated in ammonia.
Part of the way ammonia’s behavior differs from water’s is that ammonia cannot form such strong hydrogen-bonded networks. This difference explains ammonia’s lower boiling point—the molecules are more easily pulled apart when they are heated. Many of water’s subtle interactions, including that fine balance of stabilizing and flexibility-causing effects in proteins, may not be as readily possible in ammonia.
To top it off, ammonia can aggressively attack molecules in life. Ammonia, like water, dissociates into two ions in solution (NH4+ and NH2−). This solution containing NH2- binds with protons and so attacks molecules that contain them. These molecules include a vast number of complex molecules from which the life we know is assembled. This annihilative behavior makes ammonia damaging to life on Earth and in all probability very reactive to many complex molecules elsewhere. To summarize in colloquial terms, ammonia just lacks chemical subtlety.
It would be remiss if I did not point out that ammonia has some strange and noteworthy properties, though. For example, it can dissolve metals, making an eerie blue solution of metal ions and many free electrons. Free electrons are essential in life because they are the raw material for the electron transport chains that gather energy from the environment. Superficially, we might claim that a liquid that can dissolve electrons would offer a source of this most sought-after commodity. Eerie blue aliens absorbing tasty electrons from their environment in oceans of ammonia? Let us not discount it.
Despite what I have discussed so far, ammonia can be involved in complex chemistry. It is used as a solvent by industrial chemists to prepare many useful things needed by industry. It is a precursor to a smorgasbord of nitrogen-containing compounds such as hydrazine, used in rocket fuel.
The problem with ammonia, like all nonaqueous solvents proposed for life, is that we can come up with a shopping list of characteristics that are favorable for a living thing. Most liquids have some properties that do not seem inimical to life and can be, as with the solvated electrons in ammonia, possibly even useful to it. However, we are not looking for a solvent with a few things that are compatible with self-replicating, evolving organisms. We are looking for a fluid that can participate in a vast diversity of reactions and that, with the sheer breadth of chemistry we would like for building a living thing, is not too blunt-edged or reactive in its chemical behavior.
We leave the oceans of ammonia to move to other liquids that seem even less likely as useful solvents but that, in our quest to be open-minded, we might consider anyway. Some have something positive to offer scientists’ tireless quest for other liquids. Among these liquids are sulfuric acid (H2SO4), formamide (CH3NO), and hydrogen fluoride (HF).
In contrast to water, liquid sulfuric acid has a much wider temperature range, from 10 to 337°C at one atmosphere pressure, which might make it look promising since it could exist in a liquid state in a wide collection of environments. It can be found in the clouds of Venus at concentrations between 81 and 98 percent. Intriguingly, within the Venusian clouds, at around fifty kilometers high, there is a region in which temperatures are within the range of 0 to 150°C and pressures are similar to those at the surface of the Earth. The optimistic temperature and pressure data alone have invited much discussion about the possibility of life, in the shape of floating bladders bobbing along in the Venusian sky or sulfate-reducing bacteria chomping on sulfuric acid. In an intriguing thought experiment, chemist Steve Benner suggested some alternative chemistries for proteins in strange liquids. In sulfuric acid, a link between amino acids could be made stable with a sulfur atom, instead of nitrogen. Although, like water, sulfuric acid can dissolve many compounds, it is not kind
to organic material or to much complex chemistry. Its chemically destructive nature means that any biochemistry that evolved in it would likely be very limited.
Similar chemically limited vistas are found in formamide. Although many molecules, including some familiar to us, such as ATP, are stable in the substance, the smallest amount of water in combination with formamide hydrolyzes them, destroying them, meaning that oceans of formamide would have to be on an almost-waterless planet.
Not dissimilar in chemical character to water, hydrogen fluoride can form hydrogen bonds and will dissolve many small molecules. However, it is alarmingly reactive when mixed with water to form hydrofluoric acid. In the laboratory, geologists use this acid to dissolve away rocks, etching out fossils to make them easier to see. It’s propensity to react with carbon-hydrogen bonds and turn them into carbon-fluorine bonds might stymie some of its attractiveness as a solvent for organic chemistry, unless a life form can be built with a fluorine-rich set of molecular ingredients.
In addition to the challenges just discussed, other problems may come into play with the theoretical alternatives to water in our search for life-supporting fluids. This is particularly the case for those liquids that are proposed to operate at low temperatures, such as liquid ammonia.
Chemical reaction rates proceed according to a very simple principle expressed in the Arrhenius equation. Svante Arrhenius, a Swedish Nobel Prize–winning chemist and physicist, was an extraordinary polymath of the nineteenth and early twentieth centuries. He dabbled in a vast number of subjects and even speculated on the effects of adding more CO2 to the Earth’s atmosphere, predicting that it would prevent ice ages and warm the Earth. He recognized that the rate of chemical reactions depends on their temperature. Using the rates of different reactions measured in the laboratory, he showed that this dependence was not a simple linear relationship. Doubling temperature does not just increase reaction rates by the same amount regardless of temperature. The relationship is exponential. More precisely, the rate of any reaction (k) is given by
The Equations of Life Page 20