Overview;… The Origin Of Life On Earth;… Prebiotic Earth;… Life Is Water Based;… Constant Molecular Collisions;… A Mole Is A Very Specific Number Of Molecules;… Water Is Liquid Because… Acid, Alkali And pH;… Oxidation And Reduction—Electron Donor Seeks Electron Grabber And Vice Versa;… First Life Was Anaerobic;… Carbon Chemistry Is The Chemistry Of Life;… Putting Out Fires;… Life Produces And Depends Upon An Unstable Atmosphere;… The Elements of Life.
Overview
Four and a half billion years ago the cooling crust of our planet was rocked by massive earthquakes, flooded by enormous lava flows, overturned by molten iron and other heavy elements sinking toward the center and pounded incessantly by falling star trash. The hot and wet atmosphere lacked free oxygen and was subject to violent electrical storms. Three and three-fourths billion years ago our water- based solute-laden solar-powered grease-coated little ancestors left the first unmistakeable fossil traces of their organized and productive existence. This chap- ter reviews certain interactions between atoms that undoubtedly played a crucial role in the origin of life on Earth.
The Origin Of Life On Earth
Life on Earth is generally viewed as inexplicable, highly irregular, possibly even miraculous. Although a miraculous origin for life would conveniently bypass the ordinary laws of physics and chemistry, there is absolutely no evidence that any miracle has ever taken place anywhere. Nor has the most intense investigation of life’s chemistry revealed even one extraordinary reaction that violates any otherwise valid law of physics or chemistry. Indeed those seeking life’s origins frequently rep- licate life’s chemical processes in vitro (outside of the living entity) under controlled laboratory conditions in hopes of uncovering chemical events relevant to life’s start.
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Nevertheless, scientists remain a long way from understanding life, let alone con- structing it from scratch within a test tube. And while common experience assures us that it takes life to beget life, common sense also tells us that it all started some- where, somehow. Not surprisingly, there are many conflicting ideas about where and how life began. Certainly Earth’s oldest rocks show no evidence of life being present at their formation and Earth’s earliest fossil-bearing rocks carry only traces of the simplest life forms. Increasing structural complexity then becomes com- monplace among plant and animal fossils from progressively less ancient strata until the present day when Earth is obviously infested with all forms of life from the uppermost reaches of the atmosphere to many kilometers underground. Evi- dently life on Earth began very long ago and has gradually become more specialized since that time.
Some scientists attribute life’s origin to statistically improbable interactions among preexisting complex chemicals under drastic prebiotic (before life) condi- tions. This implies that even if we could exactly replicate those conditions in an unlimited number of test tubes, we still might have an indefinitely long wait before something stirred or crawled out (of course, if it was based upon almost impercep- tible chemical changes and displayed little more vigor and personality than a rock, you might not notice that something underfoot today). Others expect life to appear via numerous alternate pathways wherever energy-acquiring self-replicating struc- tures can survive. Thus they view life as an unavoidable universal process rather than an amazing outcome of some unique local event that requires very specific and convoluted explanations. And some insist that life never originated on Earth at all, that it simply saturates the entire Universe in primal spore or seed form, settling upon all possible locations as mold settles on fresh bread. Such a relegation of life’s origins to incredibly ancient times and sufficiently far away places surely fuels specu- lation—indeed, it gives full freedom to imagine the most wonderful mechanisms and miracles. But with absolutely no evidence for life’s extraterrestrial origins (or even of unearthly oversight or governance of our existence) we might as well seek life’s beginnings right here on Earth. Indeed, a lot of evidence suggests that life really is indigenous to Earth, including the way it has developed and prospered with total dependence upon locally available rather than exotic materials and processes—which also implies that exotic life forms from elsewhere would (at least initially) be ill-suited for competition with native life forms (especially the bacteria) now flourishing on Earth.
Prebiotic Earth
In view of the drastic prebiotic conditions under which First Life originated and prospered, it would likely find Earth’s present surface environment rather un- bearable. Yet any or all of the prevailing conditions on prebiotic Earth may have been essential, helpful, irrelevant or directly harmful to First Life. Perhaps only the very unusual circumstances of some isolated microenvironment deep underground made life possible at all. More likely, however, prevalent conditions mattered a lot in the assembly and survival of First Life. So what were the chemical possibilities on our enlarging Earth as it swept up nearby matter from within the rotating solar disc? Well, Earth’s continuous battering by huge numbers of gravitationally at- tracted meteorites of all sizes must have vaporized significant portions of those colliding rocks. That would have released gases such as CO, CO2, SO2, H2O, H2, N2, H2S, COS, CS2 and HCN into Earth’s growing atmosphere, while also produc- ing various hydrocarbons and reactive carbon-containing molecules including ethanol (CH3CH2OH).1 And Earth’s developing atmosphere probably included many gases such as methane (CH4) and ammonia (NH3) that are still prevalent on the outer gas planets of our solar system as well as in interstellar space. Catastrophic collisions with larger comets and meteorites undoubtedly helped to heat Earth’s surface toward its melting point (3000–3500 °K). One particularly cataclysmic impact with a planet-sized object apparently gave rise to the moon while also blow- ing away Earth’s original atmosphere—which included far greater volumes of noble (chemically inert) gases such as helium and neon. The moon’s persistent gravita- tional effect then stabilized Earth’s spin axis (obliquity with respect to the sun) and thereby prevented Earth’s climate from varying chaotically.
Our planet’s cooling surface included about 90 different elements. Some of those atoms entered readily into an endless variety of molecular arrangements based upon shared electron bonds. Others were held together ionically, often in crystal- line form, as part of Earth’s rapidly forming rocky crust. In addition to lava and rocky debris, the weathering of early continents undoubtedly produced sand, silt and clay (but no soil, since that term implies live organic content). Innumerable bodies of standing or flowing water above and under the ground must have offered our early water-based ancestors access to every possible dissolved mineral, tempera- ture and habitat. Certain of the endlessly varied and widely available clay or iron pyrite (fool’s gold) surfaces may have provided essential templates for particular molecular syntheses and replications. Energy for production of life’s complex mol- ecules came from frequent electrical storms and intense solar ultraviolet light that
1Mukhin, L. et al. 1993. “Origin of Precursors or Organic Molecules During Evaporation of Mete- orites and Mafic Terrestrial Rocks”. Nature 340. pp 46–47.
penetrated Earth’s thick clouds. Despite sunlight-induced disruption of water va- por in the upper atmosphere, the warm and wet lower atmosphere lacked free oxygen—therefore many fragile and complex molecules persisted that might have deteriorated in an oxygen-rich atmosphere. Perhaps life began almost at once. Pos- sibly it took a long while to get started. Most likely it was extinguished repeatedly before it finally could endure. Indeed, the mere fact that life has continued against all apparent odds strongly suggests that young Earth was pregnant with possibili- ties for new life.
Life Is Water-Based
Life on Earth depends upon water (H2O) and water accounts for about two- thirds of your body weight. Although the primordial protons and electrons of every hydrogen atom in the Universe were fully formed within four minutes after the Big Bang, water’s oxygen atoms only entered production following first starlight. More specifically, Earth’s oxygen atoms were individually created deep within countless overweight stars through a complex series of energy-releasing steps that allowed the strong nuclear interaction to fuse eight protons with a similar number of neutrons.
Upon being blown out of their home star, those newly formed oxygen nuclei quickly attracted a full inner shell of two electrons, while gathering just six electrons for the second (outermost or valence) shell which comfortably seats eight. Therefore every electrically neutral oxygen atom avidly seeks two additional electrons in order to complete that second and outermost electron shell. When unable to find more generous partners, oxygen atoms will join in covalently double-bonded pairs as mo- lecular oxygen (O2 or O=O)—a relationship that requires each oxygen atom to place two of its precious electrons into joint custody so both partners can complete their second shell.
Covalently-bound atoms are stuck with each other. While they may twist and vibrate a bit within the confines of this arrangement, they hesitate to break it off because of the valuable property that they share (those sparkling electrons). None- theless, few relationships are eternal, so if some fast-traveling group of hydrogen molecules happens upon a bunch of oxygen molecules lusting for more electrons, it does not take much heat (a match perhaps, or a spark of static electricity) to boost local vibrations and collisions to a point where some highly reactive atoms can break free and change partners. Such a reaction will tend to continue in a self- sustaining and markedly exothermic (heat releasing) fashion—indeed, the molecular products of this chemical reaction are so much more stable than its starting re- agents that H2 gas combines explosively with O2 gas to form water. Having thus
filled their outermost electron shells and established a more stable relationship (at a lower energy state), each oxygen atom with its two covalently bound hydrogen partners becomes quite non-reactive. This rapid and complete chemical reaction (H2 plus 1⁄2 O2 becomes H2O plus a lot of heat) explains why pilots prefer to fill their lighter-than-air dirigibles with non-reactive helium (42He) even though that costly gas weighs twice as much as hydrogen (11H2).
Constant Molecular Collisions
Atoms and molecules in the gaseous state collide continuously—therefore they strike one another or the container wall or any intruding object such as your finger with the same average impact—which means lighter atoms and smaller molecules must move far more rapidly than their more massive cousins. Further- more, average molecular velocity always rises with temperature (indeed, molecular velocity and temperature define each other) so hot air balloons go up because their enclosed air molecules hit harder and therefore occupy more space (hence weigh less per unit volume) than an equal number of cooler gas molecules outside. De- spite making up three fourths of all visible mass in the Universe, free hydrogen molecules are rarely found within Earth’s atmosphere—not only are those light- weight H2 molecules very reactive, they also travel at extremely high average velocities.
So if some nearby O2 doesn’t grab it first, several strong bumps in an outward direction are likely to kick an H2 right out of Earth’s dilute upper atmosphere—the escape velocity necessary for any molecule or spaceship to escape from Earth’s gravi- tational pull diminishes with increasing altitude, being about 1.4 times the velocity required to maintain a circular orbit at that altitude (disregarding air resistance).
From Earth’s surface that escape velocity would be about 11 kilometers per second (7 miles per second or over 25,000 miles per hour). Escape velocity from our moon’s surface is only 2.4 kilometers/sec (1.5 miles/sec)—close enough to the average speed of its atmospheric gas molecules so that the moon cannot long retain primordial gases still leaking from its interior or the other gases produced by radioactive decay of heavier atoms or during meteorite impacts.
An ordinary gas molecule at sea level in Earth’s atmosphere must endure over ten thousand collisions during the exceedingly brief moment that it takes to travel one millimeter. It is even worse in liquid water, where an H2O molecule will suffer literally millions of crashes along any one millimeter route at a staggering rate of 1013 (10 trillion) collisions per second. Such an immense number of collisions so encourages mixing and matching that nearby molecules in any water solution will surely meet and interact if suited for one another—and it only takes a picosecond
(10–12 second, or a trillionth of a second) to make or break a covalent molecular bond. When hot volcanic rock containing ferrous iron contacts sea water, each hot Fe++ will rip the oxygen heart right out of some nearby water molecule, thereby releasing its hydrogen. Solar photons similarly disrupt water molecules in the up- per atmosphere. But the customary collisional recapture of free hydrogen molecules by atmospheric oxygen has allowed Earth to retain at least a third of its original H2O allotment until now.
However, our sun is gradually becoming hotter and brighter as its helium production raises the density and thus temperature and fusion rate at the solar core. As Earth’s stratosphere (upper atmosphere) then slowly warms and becomes more humid, solar ultraviolet light above the ozone layer will increasingly photo- dissociate that water vapor—the resulting escape of free hydrogen to outer space should eliminate Earth’s remaining water within the next 2.5 billion years (a couple of billion years before our sun enters its red giant phase). Of course, Earth’s higher life forms are unlikely to survive those rising temperatures for even another billion years unless Earth can be partially shielded from that enhanced sunlight.2 But you ought not worry too much about that prediction since catastrophic cometary or asteroid impacts (plan on one every 500,000 years or so) will probably have eliminated another two thousand dominant life-forms or major civilizations by then. Furthermore, all of the above eventualities pose far less risk to you and your loved ones than your driving, your cooking, your slippery bathtub, your medicine cabinet and human overpopulation in general.
Microscopic particles of bacterial size (barely visible through an ordinary light microscope) are seen to vibrate or dance about when viewed in a water droplet.
That Brownian motion is due to cumulative variations in the random impacts of millions of water molecules from all sides. Not knowing that Brown and others had already observed this phenomenon, Einstein predicted (in 1905) that if atoms and molecules were actual physical entities rather than mere convenient fictions, an investigation of the movements of tiny particles suspended in water should confirm both the existence of molecules and their calculated size. Perrin then mea- sured those movements with such accuracy that he was able to calculate Avogadro’s Number (see below). And modern techniques such as near-field optics, scanning tunneling electron microscopy and atomic force microscopy even allow the imag- ing and manipulation of individual atoms and molecules. In any case, at identical temperatures and pressures, a specific volume of oxygen gas (O2) weighs 16 times more than an equal volume of hydrogen gas (H2)—so hydrogen is far lighter than air. More generally, equal volumes of any atomic or molecular gas at equal tem-
2Caldeira, Ken and Kasting, James F. 1992. “The Life Span of The Biosphere Revisited”.
Nature 360. pp 721–23.
peratures and pressures must contain equal numbers of independently moving at- oms or molecules. That rule, known as Avogadro’s Law, was proposed a century before Einstein’s insight and Perrin’s measurements.
A Mole Is A Very Specific Number Of Molecules
The output of any chemical reaction is limited by the number of molecules of each substance that are available to interact. This number of molecules can be determined from the weight of each substance involved but such calculations are cumbersome. It is far easier to deal in portions or multiples of the specific huge Avogadro’s number (6.02 x 1023) of molecules that defines a mole. A mole of any pure substance is one gram molecular weight—its molecular weight expressed in grams. In other words, add up the Table of Elements weights of every atom co- valently bound within a single molecule of a pure substance, weigh out exactly that many grams of that substance and you will have 6.02 x 1023 molecules or one mole of that substance. And since the gram atomic weight of He is 4, we know that 4 grams of helium contains 6.02 x 1023 atoms. Similarly, because one mole of H2 (2 grams) plus 1⁄2 mole of O2 (16 grams) can react explosively to produce one mole of H2O (18 grams), we know that 18 grams of H2O contains that same huge number of molecules.
Water Is Liquid Because…
Life depends upon water being liquid at ordinary temperatures and pressures.
Yet H2O molecules are lighter than O2 molecules (18 vs. 32 grams/mole) and O2 is a gas—so why isn’t H2O a gas? It turns out that water molecules strongly attract each other (as we might have guessed from the way a water droplet rounds up as its molecules huddle together). This self-attraction or surface tension of water comes about because every water molecule is bent (H/ O / H) rather than straight (H–O–H). In addition, each oxygen atom has seven protons more than either of its hydrogen partners, so the covalently shared electrons of a water molecule spend more time around their O than about either H. That relative excess of electrons in current possession leaves the O mid-section of every H2O somewhat negative while both H end-sections of every water molecule bear a slightly positive charge. Since both H’s of each bent water molecule can be viewed as offset to the same side of the O, the entire water molecule has one sort-of-positive side (the H’s) and a sort-of- negative side (the O). That slight separation of charges makes H2O a polar molecule which will orient itself with respect to a strong electrical field. However, being electrically neutral overall (thus not an ion), a water molecule is not pulled toward
a positive or a negative electrode in order to deposit or pick up an electron.
Now the interesting part. Put a bunch of water molecules together and every relatively negative O on one molecule will attract the relatively positive H ends of its neighboring water molecules while each of its own covalently-bound H’s are similarly drawn to some nearby relatively-negative O. Thus despite all of its bopping about, a water molecule will on average be hydrogen bonded to 3.4 closely adjacent water molecules. Although covalent bonds are about twenty times stronger than such H-bonds, the many H-bonds between water molecules should make water quite stiff (and it surely felt stiff on your last belly-flop into the pool). But fortu- nately there seems just enough room for a fifth water molecule to intermittently squeeze in between its hydrogen-bonded brethren—thereby sharing and weaken- ing one of their hydrogen bonds. So some hydrogen-bonded water molecules are always ready to rotate out of their current positions, which may explain why water flows so easily—and more rapidly under high pressure.
On the other hand, water stiffens when stretched around a hydrophobic (non- wettable) particle, due to diminished opportunities for the fifth molecule to crowd in. Similarly, each liquid water molecule moves increasingly slowly and becomes more securely hydrogen-bonded to four others during the formation of ice. Such an orderly inclusion of the fifth water molecule causes some expansion of water’s volume with freezing. That is why ice floats on liquid water while other solids (e.g.
Earth’s solid iron core) tend to shrink and settle out beneath their melted phase as they cool. Life on Earth would surely be far different if ice sank—living things might be restricted to the tropics—or maybe Earth would be a frozen lifeless ball with permanently snowy surfaces reflecting incoming sunlight back into space and little atmospheric H2O to help retain its heat (see below).
Water’s hydrogen bonds also explain why it takes so much molecular velocity (heat) to fling each water-loving H2O molecule out of liquid water as free gaseous H2O. Thus evaporation causes cooling because any water molecule that does man- age to evaporate must have been traveling far more rapidly than most of those remaining in liquid form—since temperature is defined by average molecular ve- locity, each high-speed departure of a vaporized H2O molecule reduces the average velocity of the liquid water molecules remaining. Furthermore, water molecules that contain 168O generally travel faster than those burdened with the two extra neutrons of the 188O isotope. So during Ice Ages when more of Earth’s water became tied up in snow and ice, sea water and sea life included a higher percentage of the less easily evaporated 188O isotope. Water molecules are occasionally bumped di- rectly out of ice in the process known as sublimation. Thus ice may evaporate gradually without ever passing through a visibly liquid water phase.