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http://rutgerspress.rutgers.edu/press_copyright_and_disclaimer/default.htmlPrelude
Peering through a saucer-sized porthole at the seafloor, three scientists gaped at an iron-drenched landscape. An iron-rich hot spring was bursting from the ocean's floor, precipitating towers of iron compounds, spreading iron fluff for miles around, and feeding hosts of never-before-seen primitive creatures. How did these organisms survive, even thrive, in complete darkness and under intense pressure in an environment devoid of oxygen?
Today many biologists and geologists view these iron-filled seafloor communities as living laboratories, mimics of conditions that existed when life first evolved. We know that iron formed our planet's core and permeated its mantle and crust, but did iron also play a role in life's birth? Staring through the porthole of their research submarine in 1976, scientists could only wonder at the answer.
An Iron-Rich Planet Forms
Iron's abundance on Earth is a relic of our astrophysical history. Matter in the universe began as hydrogen about 15 billion years ago. As massive stars formed and then died in supernova explosions, nuclear energy processes in the burning stars produced heavier elements, including carbon, then silicon, and finally iron.
In the midst of all this fire and brimstone, the nuclei of iron remained stable, accumulating in massive numbers. The stability of an element depends on how much energy binds together the protons and neutrons in its nucleus. Nuclei are caught between two opposing forces: one favoring a large nucleus and the other a small one. Iron stands at the equilibrium point between these forces, and the twenty-six protons and thirty neutrons of its most stable and common isotope, 56Fe, are held together more tightly than those of any other nucleus in the universe. In the heat of the early universe, elements with nuclei lighter than iron's tended to fuse together to make heavier elements, while elements with nuclei heavier than iron tended to fission apart into lighter elements. As gravity forced a platter of whirling particles to coalesce into Earth 4.6 billion years ago, it pulled the molten iron droplets toward the center, where iron alloys formed a solid inner core and a molten outer core. In this maelstrom, Earth became an iron planet.
Earth's first billion years were a hellish baptism by fire. Bombarding the planet, meteorites and asteroids incinerated its surface. Oceans boiled, lightning flashed, torrents of acid rained down, and volcanoes erupted. The atmosphere filled with dust that blocked the sun's light and with sulfur compounds that smelled like rotten eggs. At first there was little or no free oxygen and no ozone layer to absorb the sun's ultraviolet rays, and Earth's surface was sterilized by radiation and by searing temperatures. As these global forces raged, free electrons forged and broke atomic bonds, and atoms organized and reorganized themselves in changing arrays of molecules. In this dynamic sphere, iron had no peers.
Iron is not only the most abundant metal element on our planet but also one of the most versatile electron exchangers. The nuclei of all iron atoms contain twenty-six positively charged protons and between twenty-eight and thirty-two uncharged neutrons. An electrically neutral iron atom has twenty-six negatively charged electrons to counterbalance its twenty-six positively charged protons. When iron atoms participate in chemical reactions, they usually transfer one, two, or three electrons to neighboring atoms to form compounds. In compounds, iron exists primarily in two forms, or charge states, ferrous and ferric. Ferrous iron lacks two electrons because it has released them to its neighboring atoms, whereas ferric iron has released three electrons to its neighbors (see figure 1.1). With easy access to such a variety of charge states, iron atoms can take part in an enormous number of chemical reactions. Few elements are as flexible when it comes to helping make inorganic and organic molecules.
Iron and Developing Life Moving beyond iron's role in forming inorganic molecules, scientists began to believe that iron also provided some of the energy needed by developing life. To the chemist, life is the dynamic movement of electrons from atom to atom. The cellular reactions that process molecules are the vigorous signatures of life. And while cellular chemistry has many players, including proteins, enzymes, and amino acids, it is directed by the science of energy transfer, thermodynamics. Atoms donate and acquire electrons in order to find the relationship that they can maintain with the least energy. As good conductors of electricity, metals do this easily. Given the abundance of iron, carbon, nitrogen, hydrogen, and sulfur on early Earth, it is not surprising that life's earliest biochemistry was driven by these chemicals. Inevitably, metals like iron that easily grab and donate electrons become the catalysts for reactions of organic compounds made with carbon, hydrogen, nitrogen, and oxygen. For example, matrices of carbon polymer chains marked much of life's early chemistry, and iron atoms spaced along the chains could transfer electrons around the matrix with alacrity. Gliding to and fro, the electrons helped build many of biology's electronic features.
Chemists refer to the transfer of electrons as reduction and oxidation reactions, or redox reactions. Reduction is the gain of electrons. Oxidation-which does not necessarily involve oxygen at all-is the loss of electrons. These so-called redox reactions are the lifeblood of cellular functioning, and iron's facile cycling of electrons made the metal an indispensable catalyst. Sliding an electron over to another atom, ferrous iron is oxidized to ferric iron. Slipping an electron off another atom, ferric iron is reduced back to ferrous iron. The cycle can be repeated over and over. Each time iron donates an electron, it provides the energy for another chemical reaction to occur.
Before molecular oxygen invaded Earth's atmosphere, iron was plentiful in its soluble ferrous state, and ferrous iron was often called upon to donate an electron to facilitate the formation of molecules. More recently, biochemistry and genetics have clearly demonstrated that living organisms are chemical machines, progressing from simple, unicellular forms to more complex organisms. Over millions of years, nonliving molecules gave rise to living, replicating, and self-assembling cells. And as multicellular life forms evolved, their biochemical systems also grew increasingly complex.
But what fueled the first living organisms? Later in Earth's history, photosynthesis would power living cells, turning sunlight into energy for cells to grow and function. Was there, however, an energy source to fuel living organisms before photosynthesis? Could it have involved iron and sulfur? So far, no one knows; although iron's flexibility certainly makes it a contender. In the meantime, new discoveries are challenging the leading origin-of-life theory, known as the primordial soup.
For many years, this theory was embodied in the famous experiments conducted by Stanley L. Miller in 1953. Miller was inspired by the Nobel Prize-winning chemist Harold Urey and by a book, The Physical Basis of Life, written by English physicist J. D. "Sage" Bernal. Urey and Bernal suggested that the organic compounds needed for life might have formed from some of the components of Earth's early atmosphere: methane, ammonia, water, and hydrogen. To test their hypothesis, Miller simulated primordial conditions with a mixture of methane (CH4), ammonia (NH3), and hydrogen (H2). Then he ignited it with a lightning-like electric discharge. "The water in the flask became noticeably pink after the first day, and by the end of the week the solution was deep red and turbid," Miller observed. Analyzing the mixture, he discovered a rich soup of amino acids. He assumed he had produced amino acids under conditions that were likely to have existed on primitive Earth.
Miller's experimentwas quick and simple, and soon other researchers were generating many of the small organic molecules found in present-day living cells, including amino acids and sugars and the purines and pyrimidines that are components of DNA. Scientists who carried out experiments like Miller's assumed that the compounds important in biology today were the precursors of primordial life. Indeed the organic compounds found in meteorites are quite similar to the products of these experiments.
Stanley Miller's recipe for the beginning of life remains widely accepted, and many high school biology students are still taught that it is the only biochemical theory of life's origin. But many scientists believe that other, more recent and testable theories and discoveries should also be considered. Some researchers believe that Miller's recipe will have to be entirely replaced.
Miller's experiments reflected a belief that the compounds in early life must have contained carbon, hydrogen, and nitrogen with small amounts of sulfur and phosphorus thrown in. Scientists emphasize that Earth's earliest atmosphere did not contain much molecular oxygen, and new techniques for identifying trace elements in extraordinarily large bodies of matter reveal that at least twenty elements-including iron-are essential for life. Like Charles Darwin, Miller also believed that life must have begun in shallow sunlit pools at moderate temperatures. During the late 1960s, however, microbiologists Thomas Brock of the University of Wisconsin and Jim Brierly of Montana State University found microorganisms in hot springs in Yellowstone National Park. The creatures were not only surviving, they were thriving at unheard-of temperatures: 167 degrees F (75 degrees C) and higher. Until then, biology textbooks had emphasized that life on Earth depended crucially on sunlight and photosynthesis. The discovery that hot springs could support abundant life without any sunlight at all struck a powerful blow at Miller's soup recipe. What would replace it?