In this issue, your old Quantum Butcher’s going to start something new—a series on frontiers in biology. We’ll talk about the latest revelations in cell structure and physiology, recombinant DNA and genetic engineering, cell suicide, what’s new in neurobiology, cell signaling and other groovy stuff. I won’t throw it all at you sequentially—along the way I’ll be sticking in a few articles on physics, computer science, astronomy and any other topics that jump up in my face. But for the next year or so we’re going to take a journey through the frontiers of living matter—or as I like to call it, the Quantum Meat.
Some of you, I’m sure, could dig into the Quantum Meat and feel right at home munchin’ down on phospholipid bilayers, synaptic vesicles, mitochondrial membranes and retrotransposons. Let’s face it: people like you aren’t normal. Most decent, regular folk last saw a biology text in high school or college, which means, of course, that they’ve never had a proper introduction to the life sciences. So we’re going to start with fundamentals. The first three articles in this series will cover the basics, cooking up the Quantum Meat from scratch, as it were. We’re going to start with atoms and work our way up to cells. But even in these introductory articles we’re going to cover some material you bionerds may find interesting, so don’t run off to play with your PCR just yet.
Let’s wipe down the older particle cleaver and dissecting tray and get started, shall we?
FUNDAMENTALS I:
BUILDING BLOCKS—ATOMS AND BASIC CHEMISTRY FOR LIFE SCIENCES
1. Bohr’d with Atoms?Life is made of matter. And matter, for the most part, is made of atoms. There are exceptions to that latter statement, of course, but for our purposes this generalization is entirely acceptable. Atoms have undergone quite a metamorphosis since they were cooked up by the ancient Greek philosopher Democritus. In the last century or so, the atom has matured from a "plum pudding" model, in which negatively charged electrons were embedded like raisins in a positively charged goo, to the now-familiar and ubiquitous "solar system" atom, in which a massive nucleus consisting of positively charged protons and electrically neutral neutrons sits at the center and tiny electrons orbit the periphery like planets. This is more properly known as the Bohr model of the atom. The great Danish physicist Niels Bohr first formulated and used this model to attack certain fundamental problems in physics.
The Bohr model of the atom is just that—a model. And in many respects, it is a deeply flawed and incorrect model. Yet it remains an extraordinarily useful tool in physics, chemistry and biology, and we’ll be relying on it heavily. This is because the Bohr model makes it easy to see what’s going on with the formation of chemical bonds—the bonds that bind atoms together into molecules. The formation of chemical bonds is a subject at the heart of chemistry, and chemistry is at the heart of biology. So let’s take a look at some Bohr atoms.
Hydrogen and helium are the two most ancient, abundant and primitive elements in the universe. As can be seen from Figure 1, hydrogen consists of a single proton and a single electron, while helium contains two electrons surrounding a nucleus of four elementary particles—two protons and two neutrons. Now, it’s possible to add neutrons to either atom without changing the element. For example, if we added a neutron to hydrogen, it would be almost twice as heavy (since protons and neutrons have about the same mass), but it would still be hydrogen. We would call it deuterium, which is the special name for this isotope of hydrogen, but it would be hydrogen nevertheless. Add a neutron to helium and you get a heavy isotope of helium. But if you added a proton to helium, you wouldn’t have helium anymore—you’d have an isotope of lithium.
This brings up an important fundamental concept: an element is defined by the number of protons in its atoms. Not electrons. Not neutrons. The number of protons in the atom is the atomic number, and the atomic number identifies the element. Any atom with an atomic number of 2 is an atom of helium, no matter how heavy, no matter how many electrons or neutrons it has.
Neutrons have no electric charge (hence the name), but protons are positively charged and electrons are negatively charged. So look again at Figure 1: both atoms are electrically neutral. The positive charges are balanced by the negative charges. This isn’t always the case. Sometimes atoms have missing electrons or extra electrons. This results in an electrically charged atom, called an ion. You can (with considerable effort and energy) strip an electron off of helium and make a helium ion with a charge of +1, written as He+1. If you really want a challenge, try adding an electron to helium. You’ll have a whopping power bill and a very short-lived helium ion with a charge of –1: He-1.
Other atoms become ions much more readily—for example, all you have to do to make sodium ions is toss a hunk of sodium metal into water. Stand back! After the smoke clears, you’ll have a solution of sodium ions. Each of them is a sodium atom stripped of one of its electrons. Since the symbol for sodium is Na, we write: Na+1. The point is, it’s still sodium, because the atomic number—the number of protons—hasn’t changed.
Now, the atomic weight of an element is a function of how many particles it contains, primarily protons and neutrons, since the mass of the electron is relatively negligible. Look at hydrogen. Its atomic weight is about one. Helium, with two protons and two neutrons, has an atomic weight of about four. A century and a half ago, a Russian chemist named Dmitri Mendeleev demonstrated that if you write down the atomic weights of the chemical elements, you will see that similar chemical properties occur in a periodic fashion. That is, start with hydrogen, atomic weight one. It has unique chemical properties. Those chemical properties are very similar to those of lithium, atomic weight 3, and also to sodium, atomic weight 11, and potassium, atomic weight 19. Notice the interval of 8 between lithium, sodium and potassium. This is highly suggestive of a pattern correlating atomic weight with chemical properties, but of course the pattern starts out with an interval of 2, not eight. Similar imperfect patterns were apparent, in part because Mendeleev had to work with atomic weight rather than atomic number. Mendeleev got quite dour and scruffy-looking before he figured it all out. His work eventually led to the development of the periodic table of the elements.
The periodic table is one of the great achievements of the human mind. It correlates and organizes a vast body of physical knowledge. Using the periodic table, we can at a glance decide which elements are chemically related, which are heavier, which have larger atoms, and which are most unstable. Take a look at the first column in the periodic table. Here we find hydrogen, lithium, sodium, potassium, rubidium, cesium and francium. As we’ve already noted, these elements have similar chemistry, and are collectively referred to as the Group 1a metals or alkali metals. Or take a look at the elements in column 7a, which contains chlorine, fluorine, bromine, iodine and astatine—the halogens, which tend to be highly reactive. Right next to the halogens, we find the noble gases: helium, neon, argon, krypton, xenon, and radon. All of these elements are very non-reactive, even krypton which, as we all know, reacts only with Superman. This is where the noble gases get their name: Superman was the descendant of a noble family on his home world of Krypton*.
As we go down a group (column) of elements on the periodic table, we find that the atomic weights increase: potassium is more massive than sodium, xenon more massive than krypton. And as you move across the table, you find trends in atomic radius, tendency form acids or bases, etc. But the most striking feature of the table, and the most useful one for our purposes, is the chemical periodicity of the elements as represented by the various groups.
What accounts for this periodicity? Well, let’s look at two of the group 1a metals: lithium and sodium.
Inspect Figure 3. Lithium is atomic number 3, with an atomic weight of six and three electrons. Sodium is atomic number 11, with atomic weight of 22 and 11 electrons. So far we haven’t found much of a similarity. What’s more, we find that the electrons of lithium are arranged in two concentric "shells", while those of sodium are arranged in three shells. These "shells" are quantum mechanical constructs which depict the energy levels of electrons—the closer the shell is to the nucleus, the lower the energy level. Each shell can hold a specified number of electrons and no more—part of a law of physics known as the Pauli exclusion principle. Still, with two shells to lithium and three to sodium, the two elements appear as if they couldn’t be more different.
Let’s look closer. Quantum physics says that the innermost shell of an atom, the K shell, can hold no more than two electrons. Both atoms have full K shells. The next higher energy level, the L shell, can hold up to eight electrons. We note that sodium’s L shell is full, but lithium has only one electron in its L shell, and no M shell. Sodium has an M shell, with a single electron.
Am I being too obvious? By now you’ve probably realized that each atom has only one electron in its outermost shell. The same is true for all the other elements in the Ia group, from tiny hydrogen to massive Francium. And this is the key, for it is the configuration of electrons that gives an atom its chemical properties.
Why? you ask. What is it about the electrons that gives an atom its peculiar chemistry? Good question. It’s finally time to talk about chemical bonds.
2. The Ties that Bind Let’s do a little chemistry. Say we have some hydrogen gas, some chlorine gas, a tiny spark, and a big insurance policy. We can create hydrogen chloride and a most impressive explosion. Chemists write the reaction thus:
H2 + Cl2 ® 2 HCl
Okay, what does all that mean? H2 is the symbol for a molecule of hydrogen gas, which exists as two atoms of hydrogen bound to each other. Cl2 is the symbol for a molecule of chlorine gas, which exists as two atoms of chlorine bound to each other. And 2 HCl means "two molecules of HCl." A molecule of HCl, as you’ve probably deduced, is an atom of hydrogen bonded to an atom of chlorine. So this equation tells us that if we combine hydrogen gas and chlorine gas, we get hydrogen chloride gas.
But it doesn’t tell us how it happens. Let’s go back to the Bohr model.
Whoa! What a messy method of notation. And you may be shocked to learn that many chemists aren’t very, shall we say, artistically oriented. Or patient, for that matter. We need a notation that tells us more than our empirical equation, but isn’t so cluttered and tiresome as figure 4. Such a notation exists—it’s called the Lewis Structure, a simplification of the Bohr model that makes use of the octet rule.
Loosely stated, the octet rule says that the most stable configuration of an atom is a noble gas configuration. If we were to draw Bohr structures for all the noble gases, we’d find that these nonreactive (stable) gases all have eight electrons in their outermost shells (except for helium which only has one shell and fills at two electrons). So the octet rule tells us that atoms like to have eight electrons in their outermost shell. But for most atoms, this isn’t the case. Look at chlorine: it has atomic number seventeen. That means it has seventeen protons. In an electrically neutral state, it has seventeen electrons. Two electrons in the first shell plus eight in the second gives ten, which means the outermost shell has seven. Chlorine is short an electron in its outer shell. It doesn’t like that. And that’s why it comes in twos! Look at the messy Bohr structures in Figure 4, and you’ll see that the two chlorine atoms in a molecule of chlorine gas are sharing an electron. The chlorine atoms are cheating, so to speak, in their attempt to achieve a stable configuration of eight electrons in the outermost shell. By sharing an electron, the two chlorine atoms manage to (sort of) fill out their outer shells and achieve a low-energy, stable configuration.
The same goes for hydrogen, although like helium it doesn’t exactly follow the octet rule. That’s because hydrogen and helium have only one shell, and that shell achieves a noble gas configuration, an "octet" as it were, at two electrons. But hydrogen has only one electron in that single shell. So when two hydrogen atoms meet, they share their electrons to complete their outer shell and achieve a stable configuration like the noble gas helium.
We’re now ready to use Lewis structure. The Lewis structure ignores the inner shells of an atom, as they have virtually nothing to do with chemical reactivity. The Lewis structure uses dots to represent the electrons in the outer shell. A noble gas such as argon would be written as the chemical symbol, Ar, with eight dots around it. The explosive creation of hydrogen chloride from hydrogen gas and chlorine gas would be written thus:
Now it’s easy to see what’s happening. Hydrogen and chlorine combine by sharing electrons in their outermost shells, achieving a noble gas configuration for both elements. By using Lewis structures and the octet rule, we can predict with reasonable accuracy how elements will combine to make molecules and form compounds. When two elements like chlorine and hydrogen share electrons to complete their outer shells, they are said to form a covalent bond. Covalent bonds, as we shall see again and again, are absolutely essential for the construction of large bio-molecules and living systems.
So why the explosion? When chlorine and hydrogen combine to form hydrogen chloride, a great deal of energy is released. Looked at another way, if you have a volume of hydrogen chloride gas, you must expend energy to separate them and get back to chlorine gas and hydrogen gas. Your old Quantum Butcher could spend a lot of time explaining the essentials of thermochemistry, but for our purposes it’s sufficient to simply say that a chemical system tends to go from a state of relatively high energy and instability to a state of low energy and stability. The energy contained in the covalent bonds between hydrogen and chloride in the two HCl molecules is less than the total energy contained in the covalent bonds that hold together the H2 and Cl2 molecules. The difference in energy is given off as heat and light when the bonds are rearranged to give the more stable molecule. It’s important to realize that chemical bonds can store energy when formed and release energy when broken or rearranged, because chemical energy is at the heart of living systems.
Covalent bonds aren’t the only kind of chemical bonds. Let’s look at a common chemical substance, table salt. Table salt consists of chlorine and sodium, but these two elements don’t combine by sharing electrons. Rather, chlorine takes an electron from sodium, and sodium’s glad to give it up. Using Lewis structures, we write:
By now, it should be getting pretty easy for you to see what’s going on. Sodium has only one electron in its outer shell. To achieve an octet in that shell, it would have to pick up seven electrons. Now think about that. If you add an electron to sodium, it becomes a Na-1 ion. To add a second electron, you’d have to overcome the electrostatic repulsion of one negative charge for another negative charge—a lot of energy. By the time you got to the seventh electron, you’d be adding a negative charge to an atom that already had a negative charge of –6. But there’s an easier way. Sodium can just give up the electron. Now the outermost shell is the next shell in, already full at eight electrons. Sodium has attained a noble gas octet in its outermost shell and become a positive sodium ion, Na+1. Chlorine’s happy to take up that electron, because by doing so it attains an octet in its outermost shell.** It becomes a chloride ion, Cl-1. Now we have a positively charged particle (the sodium ion) and a negatively charged particle (the chloride ion). They’re very attracted to each other, and fall into an electrostatic embrace we call the ionic bond Ionic bonds are the prevailing force holding together most crystalline substances—like table salt (Figure 7).
Take a look at the periodic table. On the left hand side are metals, in dark blue, and they take up a big chunk of the table. On the far right, in light blue, are the nonmetals. We’ve now seen what happens when we combine a metal (sodium) with a nonmetal (chlorine)—an ionic bond. And this is generally the case—combining a metal with a nonmetal generally creates a salt—a substance held together by electrostatic attractions between ions. On the other hand, combining nonmetals with nonmetals usually creates compounds held together by covalent bonds. (Note that hydrogen can be treated as a metal or as a nonmetal).
There are several other types of chemical bonds, all of them weaker than ionic or covalent bonds. But the only one we really need to talk about now is a special kind of electrostatic bonding called a hydrogen bond. Hydrogen bonding, as we’ll see in future issues, is extremely important in biological systems. In order to understand hydrogen bonding, we need to finally discard our quaint illusions about subatomic particles like electrons. It’s time to look at a more sophisticated (if not more "real") version of the atom.
In quantum mechanics, particles like electrons are also waves, and their behavior can be explained in part by the Schrodinger wave equation, or "the wavefunction." On this view, the electron is nothing like a tiny speck orbiting the nucleus. In fact, there’s no way to know where the electron is at any given time, because the electron is really just a "cloud" of probability—a fuzzy region where the electron might be. If you hold a hydrogen atom in your hand, there’s a pretty good probability that the electron is located at about .40 Angstroms from the nucleus. There’s also a very small but nevertheless very real probability that the electron is somewhere on Mars. The picture of a hydrogen atom that we get from quantum mechanics is of a slightly fuzzy sphere of probability for the position of the proton surrounded by a considerably more fuzzy sphere of probability for the position of the electron (Figure 8).
Let’s picture such a hydrogen atom covalently bound to a larger molecule. We’ll call the rest of the molecule "R". It’s a big molecule, and it has a lot of hydrogens on it, but we’re only concerned with this particular hydrogen, which is bound to an atom hungry for electrons, an electrophilic atom. Oxygen is such an atom—it has a lot more protons in its nucleus than hydrogen, and it’s missing two electrons in its outer shell, so it loves to "share" electrons with hydrogen—much as the IRS loves to "share" the money you make at your job. When a bond forms between hydrogen and an electrophile like oxygen, we find that the electron cloud for hydrogen is pulled in the direction of the electrophilic atom, as shown in Figure 9. With the electron cloud pulled to one side, the proton in the nucleus of the hydrogen atom is left relatively "bare." So the hydrogen atom has become a little positive charge sticking off the end of the molecule—not a full positive charge, but a fractional positive charge we denote d +. If another molecule nearby has a negatively charged atom or group of atoms hanging off of it, the d + charge will be attracted to it, and a weak electrostatic bond will form between the two molecules. This is extremely important in living systems, because such a weak bond can be formed, broken and reformed with a minimal expenditure of energy, thereby leading to changes in the spatial relationships between molecules—or even between groups of atoms within a single molecule—with significant biological consequences.
3. Polar Molecules—What’s So Great About Water?With our discussion of hydrogen bonding, we’ve stumbled onto concepts that permit us to discuss the structure of water and its importance in biological systems.
Water, as all terrestrial living things know, is great stuff. It’s a unique molecule—it occupies three physical phases (gas, liquid, solid) at terrestrial temperatures and pressures. It’s our most abundant greenhouse gas, and therefore important for climactic regulation. It exhibits surface tension and capillary action, which allow for the efficient sequestration and mobilization of water by ecosystems and organisms. It’s less dense as a solid than as a liquid, which is very important ecologically—if ice sank, rivers, lakes and ponds would become uninhabitable in winter –in fact, they’d freeze completely. But for the cell, the most important features of water have to do with its solvent properties, and these in turn depend on its nature as a polar molecule.
What do I mean by "polar?" I mean that the electric charge in a molecule of water is distributed unevenly. Water, as we all know, is two atoms of hydrogen bound to a single atom of oxygen. In exactly the same fashion as we discussed earlier, oxygen gets an unfair share of the electron clouds from the two hydrogens. The result is a little molecule that bears a slight d - on one end and a slight d + on the other, as shown in Figure 10. Compare this situation to ethane, an organic molecule also pictured in Figure 10. Ethane is two carbon atoms bonded to each other, with each carbon carrying three hydrogens. The carbon atoms are not especially electrophilic, and the molecule is symmetrical anyway, so the electron density is evenly distributed over the molecule. Ethane, unlike water, is nonpolar.
The polar nature of water has immense physical, chemical and biological consequences, not the least of which is water’s impressive activity as a solvent. For example, water is an excellent solvent for ionic substances, like sodium chloride. As shown in Figure 11, water dissolves salt by forming molecular "cages" around sodium and chloride ions, using its negative ends to hold sodium ions and its positive ends to hold chloride ions.
But water can also dissolve some nonpolar molecules, such as ethane, by forming cages of water molecules held together by hydrogen bonds. If the concentration of such a nonpolar molecule gets too high, the nonpolar guys will all get together in a sort of molecular wagon train circle to minimize their exposure to the water—but that’s important to living systems too, as we’ll see in our next exciting episode. The point is that water is an extraordinarily versatile solvent because of its molecular structure.
What’s more, water is abundant in the universe. This is a key feature of substances found in living systems, and that’s why scientists get so flushed and trembly when they find a world with water on it, like Europa or Callisto. There are other thallasogens (sea-forming substances, a word coined by Asimov) that could conceivably support life elsewhere in the universe, but none of them combine the cosmic abundance and unique physical properties that would make them as desirable as good old H2O. Ammonia, for example, is an excellent polar solvent, and quite abundant, but it’s only liquid at very low temperatures. And when it freezes, it sinks.***
The polar nature of water and the way different solutes respond to it has a great deal to do with the shapes of cellular molecules and cellular structures. As we’ll see next time, many large molecules, like proteins and phospholipids, have both polar and nonpolar regions. These large, complex molecules will twist, bend or aggregate in such a way as to maintain conformations that minimizes the contact between water and nonpolar regions.
Read that sentence again. That’s a big chunk of biochemistry and biophysics, right there. And it all comes down to the structure and behavior of water.
Water has another great property—it can be ionized. Remember when we tossed sodium into water? We ended up with sodium ions in solution. But we also ended up with "water ions," as it were. When the sodium ionized, it did so by giving up the spare electron in its outer shell—we’ve seen this before. But that electron had to go somewhere. It went to a water molecule, where it broke the covalent bond between hydrogen and oxygen, by completing oxygen’s outer shell. Simply put, oxygen no longer needs a hydrogen to complete its outer shell. So we end up with a fragment of a water molecule called a hydroxyl ion, written OH-. It’s an ion because it’s got an extra electron and so it’s negatively charged. Note that we also end up with a hydrogen atom. But hydrogen doesn’t like to be all by itself—he’s got an outer shell to fill too, you know. So he binds to another hydrogen atom from another water molecule stripped by sodium, forming hydrogen gas. This forces us to balance the books by multiplying the entire reaction by two—for chemistry cannot create matter, and we must end up with the same atoms we started with, albeit in different combinations.. The point is this: we’ve created a new species, the hydroxyl ion, which is a highly reactive molecule and which we call a base. Throwing sodium metal into water results in a basic solution of sodium hydroxide.
Now let’s throw some of that hydrogen chloride we made earlier into water. What happens? The chlorine runs off with all the electrons, that’s what—after stealing them from hydrogen to fill out the outer shell, it’s not about to let them go now. We end up with a chloride ion and a hydrogen ion. The hydrogen ion, in reality, binds to the water to form a water molecule with an extra hydrogen—the hydronium ion. This hydronium ion, though, quickly breaks up into water and hydrogen ion, only to reform again and again. For all practical purposes, we have water, chloride and a bunch of homeless hydrogen ions, so the reaction is usually written as in Figure13(b).
Such a solution of hydrogen chloride in water is called hydrochloric acid. The acid part is the extra proton, donated by the HCl. There are many such hydrogen donors—molecules that donate hydrogen to ionize water. Hydrogen donors are acids.
Isn’t that easy? A substance, like sodium, that depletes hydrogen from water and forms hydroxyl ions is a base. A substance, like HCl, that adds hydrogen to water and forms hydronium ions is an acid. It’s that simple.
We can measure the concentration of hydrogen ions in a solution (and thereby measure the acidity). This measurement is called the pH. Technically, the pH is the inverse of the log of the hydrogen ion concentration—what this means, quite simply, is that the lower the pH, the more hydrogen ions there are in solution, and the more acid it is. The higher the pH, the less hydrogen ions there are—meaning more hydroxyl ions. A high pH indicates more hydroxyl ions and a basic solution.
Many biological molecules are acids or bases. Many are both acids and bases—as we shall see.
4. The Elements of Life Wow. We’ve covered a lot of ground. We’re in the home stretch now, but you need to bear with the old Quantum Butcher a while longer. After all, if you don’t eat your meat, you can’t have any pudding.
Now that we’ve surveyed the basics of chemistry (in about 6000 words—I think we’re all to be commended) we’re ready to talk about the elements of life. As I’ve already hinted, the elements of life are bound to have certain features—for one thing, they must be abundant. The ten most abundant elements in the universe, in descending order, are hydrogen, helium, oxygen, neon, nitrogen, carbon, silicon, magnesium, iron, and sulfur. Of these, seven are important in living systems—only helium and neon (non-reactive noble gases) and silicon are left out. The five most abundant elements in living systems are hydrogen, oxygen, carbon, nitrogen and calcium. Only calcium is out of the Cosmic Ten. But calcium is abundant in the earth’s crust, ranking eighth. The lesson for you sf writers out there is that it’s not a good idea to design your alien life forms out of large amounts of, say, chromium or lawrencium. Life is made of stuff that’s easy to get.
The other thing it’s important to notice about the top five elements of life is that only one, calcium, is truly and unambiguously a metal. Metals are important in living systems as trace elements, but the bulk of living systems is made up of nonmetals. And as we’ve already seen, that must mean that living matter is dominated by the covalent bond. This is why any vision of "crystalline" life is so alien—not impossible, mind you, just very, very weird. Covalent bonding works well for life. Unlike the ionic bonds that dominate in crystals, the covalent bond is not geometrically rigid—it can swivel. This means that biomolecules built on covalent bonds can be highly mobile and malleable. Furthermore, covalent bonds don’t dissolve like ionic bonds do. And covalent bonds store chemical energy, a feature that all living systems exploit. Much of the solar energy harvested by and channeled through the ecosphere is used to make substances full of energy-rich covalent bonds. There’s a word for such substances: food.
The elements of life combine in virtually endless ways to build the living matter I call Quantum Meat. But there are two covalent bonds which, I would argue, are more important than all the others. One of these we’ve already seen: the covalent bond between hydrogen and oxygen found in water. The other is the covalent bond between carbon and carbon.
5. Organic Chemistry Ain’t Biochemistry—But it Helps Carbon is the miracle atom for life. It has four unpaired electrons in its outermost shell, meaning it can form four covalent bonds. What’s more, it’s neither too stingy nor too generous with its electrons—carbon-carbon bonds have a range of energies that enable them to be formed and broken at terrestrial temperatures and pressures. Perhaps most important, carbon loves to bond to itself, forming chains, spirals, spheres, lattices—an endless variety of shapes. It’s the ultimate atomic tinkertoy.
We’ve seen that the quantum electron fuzz around the hydrogen atom is spherical, but the outer electrons of carbon occupy a much more complex shape, forming a tetrahedron. Consequently, the compounds of carbon branch out into three-dimensional space.
The study of the chemistry of carbon is a discipline in itself—organic chemistry. Compounds containing carbon-carbon bonds are organic compounds. Not all organic chemistry is biochemistry, but virtually all biochemistry is organic chemistry. Carbon compounds incorporate nitrogen, oxygen, hydrogen, phosphorus and sulfur to form a vast catalogue of molecules necessary for life, including sugars, phospholipids, ketones, acids, and alkaloids, among others. We’ll talk more about these molecules in the next installment. For now, just to give you a taste of carbon’s diversity, spend a little time with the crash course in organic chemistry offered in Figures 15-17.
6. Why Carbon?Recently a scientist friend of mine read a story I’d written about the exploration of Europa’s biosphere. He liked the story, but he thought he’d found a huge scientific flaw. He was dismayed that I’d based my Europan lifeforms on carbon—worse, that I’d built them with amino acids, phospholipids and other molecules found in Earthly organisms. Wouldn’t alien creatures have a completely, well, alien biochemistry? Wouldn’t they most likely be based on a different chemistry altogether, eschewing carbon for some other element with similar chemical properties?
Possibly. But not probably. Now, it’s not unreasonable to speculate about such things. Remember the periodic law—elements in a particular column of the table will have similar chemical properties. If you refer to the table, you’ll see that silicon, Si, lives right under carbon. Silicon, like carbon, can form chains and lattices, and it binds to the same family of elements as carbon does. And silicon, as we’ve already seen, is one of the ten most abundant elements in the universe. So why not base life on silicon, instead of carbon?
Let’s start with abundance. Carbon is ten times more abundant than silicon in our solar system. This isn’t to say that silicon isn’t abundant—in fact, there’s more of it in Earth’s crust than there is carbon. And that’s part of the problem. Silicon, while more abundant in the crust, is less accessible than carbon, because it’s tied up in the form of silicates and other rocky materials. What’s more, because carbon is the winner in cosmic abundance, it’s more available to cosmic processes, many of which are now known to create prebiotic molecules. In the clouds of gas and dust between the stars, many organic compounds are formed, including cyanide, carbon monoxide, and even amino acids and benzenes. Of the 75 or so molecules found in interstellar space, only ten are made with silicon.
Another extremely limiting factor for silicon as a biotrophic substance is that it cannot easily take part in biogeochemical cycles. Such cycles are increasingly seen as important to life on earth; indeed, the complex, interconnected cycles of elements on earth are regarded by many daring scientists as constituting a "physiology of Gaia." Look at carbon: it occupies huge reservoirs in the crust (as carbonates and dead organic matter), in the oceans (as dissolved CO2) in living systems (vast stores of carbon in forests, grasslands, plankton and other ecosystems) and in the atmosphere (as CO2). Carbon is continuously cycled to and from these vast reservoirs, primarily as CO2 expired by animals and inspired by plants and released from the oceans. SiO2, silicon dioxide, is also abundant on the earth—as sand. Sand, as you might imagine, doesn’t cycle much. So silicon can’t participate in the biogeochemical cycles that form the lifeblood of large interlocked ecologies, at least not at terrestrial temperatures and pressures.
Another problem is the chemical bond. As we’ve seen, silicon on earth tends to exist as silicon dioxide. Why? Because the silicon-oxygen bond is strong and stable. In fact, the silicon-oxygen bond takes twice as much energy to break as a silicon-silicon bond. This means that when you have silicon and oxygen together on a planet, you’ll tend to form silicates (silicon-oxygen compounds), not "silorganics." That’s going to significantly decrease the practical abundance of silicon for use by living systems. Now, it is possible to form silicon-oxygen polymers, but again, the bond energy of these polymers is very high, making them much more unreactive and stable than the carbon-carbon bonds in organic molecules. So silicon-oxygen polymers are relatively inert, rendering them much less suitable for use in living systems.
For these and other reasons, no other element comes even close to carbon for abundance, availability, potential for biogeochemical cycling, and appropriate bond energy and diversity. This isn’t to discourage you sf-writing-types (and I know you’re out there) from cooking up the Quantum Meat from elements other than carbon. I’m just letting you know that it’ll be difficult, and that you’ll have to make allowances in your recipe, or those of us with a discerning palate are likely to choke on your efforts.
Next time, we’ll take the basic principles of chemistry and use them to build biological molecules and basic cellular components. Bring your appetite.
In Memory of Lou Csontos
* Just kidding. Everybody knows Superman was the illegitimate son of a bored farmgirl and a travelling neutronium salesman.
** Actually, that’s a little misleading. Ionizing sodium and chlorine requires energy, but the investment in energy is paid off in full when one considers the total energy released by the formation of the ionic bond.
*** Something else to think about—life developing on ammonia-rich worlds will have a hard time developing technology, because rocks and ores will be difficult to mine and because in such an atmosphere nothing will burn.
BIBLIOGRAPHY If you want to know more, check out the following selections:
The Period Kingdom, by P.W. Atkins. Harper Collins, 1995. A fascinating perspective on the periodic table, written for the layman. Atkins looks elements as forming an actual landscape, full of peculiar yet rational topographies. An excellent introduction to general chemistry.
Gaia’s Body: Toward a Physiology of the Earth, by Tyler Volk. Springer-Verlag, 1998. A look at the biogeochemistry of earth, with emphasis on cycling, energy transformation, and protocols for "Gaian inquiry." Anybody who thinks the Gaia concept is feelgood New Age fluff had better take a look at this penetrating and hardheaded view of life on earth.
World Building, by Stephen L Gillett. Writer’s Digest Books, 1996. Of all the books in the WD Science Fiction Writing Series, this is the one that’s probably most worth your while. There’s a ton of useful information here, not only for understanding and constructing biological systems, but for constructing entire planets. Get it.
The Cartoon Guide to Genetics, by Larry Gonick and Mark Wheelis. Barnes and Noble 1983. For those of you who like to read ahead, this great little book will get you started on basic biochemistry, molecular biology and recombinant DNA engineering. Must reading.
[ ISSUE #01 ] Immortality: A Primer
[ ISSUE #02 ] Nano, Right Under Our Noses
[ ISSUE #03 ] The Keen-O Neutrino
[ ISSUE #04 ] The Keen-O Neutrino Part II, and a Little Lost Lambda
[ ISSUE #05 ] Grabbing a Computer by the Tail [ special Y2K feature by guest columnist Andrew Burt ]
Jonathon Sullivan MD, PhD, is Assistant Professor of Emergency Medicine at Wayne State University/Detroit Receiving Hospital. He is the recipient of an NIH research grant, to investigate molecular mechanisms of brain death after cardiac arrest. Dr. Sullivan lives in Farmington Hills, MI.
