第一章 绪论 生命的分子基础

Biochemistry describes in molecular terms the structures, mechanisms, and chemical processes shared by all organisms and provides organizing principles that underlie life in all its diverse forms. Although biochemistry provides important insights and practical applications in medicine, agriculture, nutrition, and industry, its ultimate concern is with the wonder of life itself.

In this introductory chapter we give an overview of the cellular, chemical, physical, and genetic backgrounds of biochemistry and the overarching principle of evolution — how life emerged and evolved into the diversity of organisms we see today.

In each chapter of this book, we organize our discussion around central principles or issues in biochemistry. In this chapter, we consider the features that define a living organism, and we develop these principles: Cells are the fundamental unit of life. Although they vary in complexity and can be highly specialized for their environment or function within a multicellular organism, they share remarkable similarities.

Cells use a relatively small set of carbon-based metabolites to create polymeric machines, supramolecular structures, and information repositories. The chemical structure of these components defines their cellular function. The collection of molecules carries out a program, the end result of which is reproduction of the program and selfperpetuation of that collection of molecules — in short, life. Living organisms exist in a dynamic steady state, never at equilibrium with their surroundings. Following the laws of thermodynamics, living organisms extract energy from their surroundings and employ it to maintain homeostasis and do useful work. Essentially all of the energy obtained by a cell comes from the flow of electrons, driven by sunlight or by metabolic redox reactions.

Cells have the capacity for precise self-replication and self-assembly using chemical information stored in the genome. A single bacterial cell placed in a sterile nutrient medium can give rise to a billion identical “daughter” cells in 24 hours. Each cell is a faithful copy of the original, its construction directed entirely by information contained in the genetic material of the original cell. On a larger scale, the progeny of vertebrate animals share a striking resemblance to their parents, also the result of their inheritance of parental genes.

Biochemistry aims to explain biological form and function in chemical terms. During the first half of the twentieth century, parallel biochemical investigations of glucose breakdown in yeast and in animal muscle cells revealed remarkable chemical similarities between these two apparently very different cell types; for example, the breakdown of glucose in yeast and in muscle cells involved the same 10 chemical intermediates and the same 10 enzymes. Subsequent studies of many other biochemical processes in many different organisms have confirmed the generality of this observation, neatly summarized in 1954 by the biochemist Jacques Monod: “What is true of E. coli is true of the elephant.”

The unity and diversity of organisms become apparent even at the cellular level. The smallest organisms consist of single cells and are microscopic. Larger, multicellular organisms contain many different types of cells, which vary in size, shape, and specialized function. Despite these obvious differences, all cells of the simplest and most complex organisms share certain fundamental properties, which can be seen at the biochemical level.

Cells of all kinds share certain structural features (Fig. 1-1). The plasma membrane defines the periphery of the cell, separating its contents from the surroundings. It is composed of lipid and protein molecules that form a thin, tough, pliable, hydrophobic barrier around the cell. The membrane is a barrier to the free passage of inorganic ions and most other charged or polar molecules. Transport proteins in the plasma membrane allow the passage of certain ions and molecules, receptor proteins transmit signals into the cell, and membrane enzymes participate in some reaction pathways. Because the individual lipids and proteins of the plasma membrane are not covalently linked, the entire structure is remarkably flexible, allowing changes in the shape and size of the cell.

All cells have, for at least some part of their life, either a nucleoid or a nucleus, in which the genome — the complete set of genes, composed of DNA (deoxyribonucleic acid) — is replicated and stored, with its associated proteins. The nucleoid, in bacteria and archaea, is not separated from the cytoplasm by a membrane; the nucleus, in eukaryotes, is enclosed within a double membrane, the nuclear envelope. Cells with nuclear envelopes make up the large domain Eukarya (Greek eu, “true,” and karyon, “nucleus”). Microorganisms without nuclear membranes, formerly grouped together as prokaryotes (Greek pro, “before”), are now recognized as comprising two very distinct groups: the domains Bacteria and Archaea, described below.

The upper limit of cell size is probably set by the rate of transport of nutrients into the cell and waste products out. As the size of a cell increases, its surface-to-volume ratio decreases. For a spherical cell, the surface area is a function of the square of the radius (r2), whereas its volume is a function of r3. A bacterial cell the size of Eschericia coli is so small, and the ratio of its surface area to its volume is so large, that every part of its cytoplasm is easily reached by nutrients moving across the membrane and into the cell. With increasing cell size, surface-to-volume ratio decreases, until metabolism consumes nutrients faster than transmembrane carriers can supply them.

In one interpretation of sequence similarities, all living organisms fall into one of three large groups (domains) that define three branches of the evolutionary tree of life originating from a common progenitor (Fig. 1-3). Two large groups of single-celled microorganisms can be distinguished on genetic and biochemical grounds: Bacteria and Archaea. Bacteria inhabit soils, surface waters, and the tissues of other living or decaying organisms. Many of the Archaea, recognized as a distinct domain by the microbiologist Carl Woese in the 1980s, inhabit extreme environments — salt lakes, hot springs, highly acidic bogs, and the ocean depths. The available evidence suggests that the Archaea and Bacteria diverged early in evolution. All eukaryotic organisms, which make up the third domain, Eukarya, evolved from the same branch that gave rise to the Archaea; eukaryotes are therefore more closely related to archaea than to bacteria.

Macromolecules and their monomeric subunits differ greatly in size. An alanine molecule is less than 0.5 nm long. A molecule of hemoglobin, the oxygen-carrying protein of erythrocytes (red blood cells), consists of nearly 600 amino acid subunits in four long chains, folded into globular shapes and associated in a structure 5.5 nm in diameter. In turn, proteins are much smaller than ribosomes (about 20 nm in diameter), which are much smaller than organelles such as mitochondria, typically 1 μm in diameter. It is a long jump from simple biomolecules to cellular structures that can be seen with the light microscope. Figure 1-9 illustrates the structural hierarchy in cellular organization.

Fewer than 30 of the more than 90 naturally occurring chemical elements are known to be essential to organisms. Most of the elements in living matter have a relatively low atomic number; only three have an atomic number above that of selenium, 34 (Fig. 1-11). The four most abundant elements in living organisms, in terms of percentage of total number of atoms, are hydrogen, oxygen, nitrogen, and carbon, which together make up more than 99% of the mass of most cells. They are the lightest elements capable of efficiently forming one, two, three, and four bonds, respectively; in general, the lightest elements form the strongest bonds.

Of greatest significance in biology is the ability of carbon atoms to form very stable single bonds with up to four other carbon atoms. Two carbon atoms also can share two (or three) electron pairs, thus forming double (or triple) bonds.

The four single bonds that can be formed by a carbon atom project from the nucleus to the four apices of a tetrahedron (Fig. 1-13), with an angle of about 109.5° between any two bonds and an average bond length of 0.154 nm. There is free rotation around each single bond, unless very large or highly charged groups are attached to both carbon atoms, in which case rotation may be restricted. A double bond is shorter (about 0.134 nm) and rigid, and it allows only limited rotation about its axis.

Most biomolecules can be regarded as derivatives of hydrocarbons, with hydrogen atoms replaced by a variety of functional groups that confer specific chemical properties on the molecule, forming various families of organic compounds. Typical of these are alcohols, which have one or more hydroxyl groups; amines, with amino groups; aldehydes and ketones, with carbonyl groups; and carboxylic acids, with carboxyl groups (Fig. 1-14). Many biomolecules are polyfunctional, containing two or more types of functional groups (Fig. 1-15), each with its own chemical characteristics and reactions.

Dissolved in the aqueous phase (cytosol) of all cells is a collection of perhaps several thousand different small organic molecules (Mr~100 to ~500), with intracellular concentrations ranging from nanomolar to > 10 mM (see Fig. 13-31). (See Box 1-1 for an explanation of the various ways of referring to molecular weight.) These are the central metabolites in the major pathways occurring in nearly every cell — the metabolites and pathways that have been conserved throughout the course of evolution.

Many biological molecules are macromolecules, polymers with molecular weights above ~5,000 that are assembled from relatively simple precursors (Fig. 1-16). Shorter polymers are called oligomers (Greek oligos, “few”). Proteins, nucleic acids, and polysaccharides are macromolecules composed of monomers with molecular weights of 500 or less. Synthesis of macromolecules is a major energy-consuming activity of cells. Macromolecules themselves may be further assembled into supramolecular complexes, forming functional units such as ribosomes.

The nucleic acids, DNA and RNA, are polymers of nucleotides. They store and transmit genetic information, and some RNA molecules have structural and catalytic roles in supramolecular complexes. The genome is the entire sequence of a cell’s DNA (or in the case of RNA viruses, its RNA), and genomics is the characterization of the structure, function, evolution, and mapping of genomes. The polysaccharides, polymers of simple sugars such as glucose, have three major functions: as energy-rich fuel stores, as rigid structural components of cell walls (in plants and bacteria), and as extracellular recognition elements that bind to proteins on other cells. Shorter polymers of sugars (oligosaccharides) attached to proteins or lipids at the cell surface serve as specific cellular signals. A cell’s glycome is its entire complement of carbohydrate-containing molecules. The lipids, water-insoluble hydrocarbon derivatives, serve as structural components of membranes, energy-rich fuel stores, pigments, and intracellular signals. The lipid-containing molecules in a cell constitute its lipidome.

The covalent bonds and functional groups of a biomolecule are, of course, central to its function, but so also is the arrangement of the molecule’s constituent atoms in three-dimensional space — its stereochemistry. Carbon-containing compounds commonly exist as stereoisomers, molecules with the same chemical bonds and same chemical formula but different configuration, the fixed spatial arrangement of atoms. Interactions between biomolecules are typically stereospecific, requiring specific configurations in the interacting molecules.

In space-filling models, the radius of each “atom” is proportional to its van der Waals radius, and the contours of the model define the space occupied by the molecule (the volume of space from which atoms of other molecules are excluded).

Configuration is conferred by the presence of either (1) double bonds, around which there is little or no freedom of rotation, or (2) chiral centers, around which substituent groups are arranged in a specific orientation.

Figure 1-18a shows the configurations of maleic acid and its isomer, fumaric acid. These compounds are geometric isomers, or cis-trans isomers; they differ in the arrangement of their substituent groups with respect to the nonrotating double bond (Latin cis, “on this side” — groups on the same side of the double bond; trans, “across” — groups on opposite sides). Maleic acid (maleate at the neutral pH of cytoplasm) is the cis isomer, and fumaric acid (fumarate) is the trans isomer; each is a well-defined compound that can be separated from the other, and each has its own unique chemical properties. A binding site (on an enzyme, for example) that is complementary to one of these molecules would not be complementary to the other, which explains why the two compounds have distinct biological roles despite their similar chemical makeup.

A molecule with only one chiral carbon can have two stereoisomers; when two or more (n) chiral carbons are present, there can be 2n stereoisomers. Stereoisomers that are mirror images of each other are called enantiomers (Fig. 1-19). Pairs of stereoisomers that are not mirror images of each other are called diastereomers (Fig. 1-20).

As the biologist, microbiologist, and chemist Louis Pasteur first observed in 1843 (Box 1-2), enantiomers have nearly identical chemical reactivities but differ in a characteristic physical property: optical activity. In separate solutions, two enantiomers rotate the plane of plane-polarized light in opposite directions, but an equimolar solution of the two enantiomers (a racemic mixture) shows no optical rotation. Compounds without chiral centers do not rotate the plane of plane-polarized light.

In this way, each chiral carbon is designated either (R) or (S), and the inclusion of these designations in the name of the compound provides an unambiguous description of the stereochemistry at each chiral center. Another naming system for stereoisomers, the D and L system, is described in Chapter 3. A molecule with a single chiral center can be named unambiguously by either system, as shown here. The two naming systems are based on different criteria, so no general correlation can be made between, say, the L isomer and the (S) isomer seen in this example.

Distinct from configuration is molecular conformation, the spatial arrangement of substituent groups that, without breaking any bonds, are free to assume different positions in space because of the freedom of rotation about single bonds. In the simple hydrocarbon ethane, for example, there is nearly complete freedom of rotation around the C— C bond. Many different, interconvertible conformations of ethane are possible, depending on the degree of rotation (Fig. 1-21). Two conformations are of special interest: the staggered, which is more stable than all others and thus predominates, and the eclipsed, which is the least stable. We cannot isolate either of these conformational forms, because they are freely interconvertible. However, when one or more of the hydrogen atoms on each carbon is replaced by a functional group that is either very large or electrically charged, freedom of rotation around the C— C bond is hindered. This limits the number of stable conformations of the ethane derivative.

The study of biomolecular stereochemistry, with precise physical methods, is an important part of modern research on cell structure and biochemical function.

In living organisms, chiral molecules are usually present in only one of their chiral forms. For example, the amino acids in proteins occur only as their L isomers; glucose occurs only as its D isomer. (The conventions for naming stereoisomers of the amino acids are described in Chapter 3; those for sugars, in Chapter 7. The RS system, described above, is the most useful for some biomolecules.) In contrast, when a compound with an asymmetric carbon atom is chemically synthesized in the laboratory, the reaction usually produces both possible chiral forms: a mixture of the D and L forms, for example.

Living cells produce only one chiral form of a biomolecule because the enzymes that synthesize that molecule are also chiral. Stereospecificity, the ability to distinguish between stereoisomers, is a property of enzymes and other proteins and a characteristic feature of biochemical interactions.

Molecular configuration can be changed only by breaking and re-forming covalent bonds. For a carbon atom with four different substituents (a chiral carbon), the substituent groups can be arranged in two different ways, generating stereoisomers with distinct properties. Only one stereoisomer is biologically active. Molecular conformation is the position of atoms in space that can be changed by rotation about single bonds, without covalent bonds being broken.

The monomeric subunits of proteins, nucleic acids, and polysaccharides are joined by covalent bonds. In supramolecular complexes, however, macromolecules are held together largely by noncovalent interactions — much weaker, individually, than covalent bonds. Among these noncovalent interactions are hydrogen bonds; ionic interactions (between charged groups); and aggregations of nonpolar groups in aqueous solution, brought about by van der Waals interactions (also called London forces) and by the hydrophobic effect — all of which have energies much smaller than those of covalent bonds. (These noncovalent interactions are described in Chapter 2.) The large numbers of weak interactions between macromolecules in supramolecular complexes stabilize these assemblies, producing their unique structures.

Polar biomolecules dissolve readily in water because they can replace water-water interactions with energetically favorable water-solute interactions. In contrast, nonpolar biomolecules are poorly soluble in water because they interfere with water-water interactions but are unable to form water-solute interactions. In aqueous solutions, nonpolar molecules tend to cluster together. Hydrogen bonds and ionic, hydrophobic (from the Greek, meaning “water-fearing”), and van der Waals interactions are individually weak, but collectively they have a very significant influence on the three-dimensional structures of proteins, nucleic acids, polysaccharides, and membrane lipids.

The result of this unequal electron sharing is two electric dipoles in the water molecule, one along each of the H—O bonds; each hydrogen atom bears a partial positive charge (δ+), and the oxygen atom bears a partial negative charge equal in magnitude to the sum of the two partial positives (2δ−). As a result, there is an electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another (Fig. 2-1b), called a hydrogen bond. Throughout this book, we represent hydrogen bonds with three parallel blue lines, as in Figure 2-1b.

Hydrogen bonds are relatively weak. Those in liquid water have a bond dissociation energy (the energy required to break a bond) of about 23 kJ/mol, compared with 470 kJ/mol for the covalent O—H bond in water or 350 kJ/mol for a covalent C—C bond. The hydrogen bond is about 10% covalent, due to overlaps in the bonding orbitals, and about 90% electrostatic.

Hydrogen bonds are not unique to water. They readily form between an electronegative atom (the hydrogen acceptor, usually oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom (the hydrogen donor) in the same or another molecule (Fig. 2-3). Hydrogen atoms covalently bonded to carbon atoms do not participate in hydrogen bonding, because carbon is only slightly more electronegative than hydrogen and thus the C—H bond is only very weakly polar.

Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction, which occurs when the hydrogen atom and the two atoms that share it are in a straight line—that is, when the acceptor atom is in line with the covalent bond between the donor atom and H (Fig. 2-5). This arrangement puts the positive charge of the hydrogen ion directly between the two partial negative charges. Hydrogen bonds are thus highly directional and capable of holding two hydrogen-bonded molecules or groups in a specific geometric arrangement.

Water dissolves salts such as NaCl by hydrating and stabilizing the Na+ and Cl− ions, weakening the electrostatic interactions between them and thus counteracting their tendency to associate in a crystalline lattice (Fig. 2-6). Water also readily dissolves charged biomolecules, including compounds with functional groups such as ionized carboxylic acids (—COO−), protonated amines (—NH+3), and phosphate esters or anhydrides. Water replaces the solute-solute hydrogen bonds linking these biomolecules to each other with solute-water hydrogen bonds, thus screening the electrostatic interactions between solute molecules.

Ionic interactions between dissolved ions are much stronger in less polar environments, because there is less screening of charges by the nonpolar solvent. Water is effective in screening the electrostatic interactions between dissolved ions because it has a high dielectric constant, a physical property that reflects the number of dipoles in a solvent.

When water is mixed with benzene or hexane, two phases form; neither liquid is soluble in the other. Nonpolar compounds such as benzene and hexane are hydrophobic—they are unable to undergo energetically favorable interactions with water molecules, and they interfere with the hydrogen bonding among water molecules. All molecules or ions in aqueous solution interfere with the hydrogen bonding of some water molecules in their immediate vicinity, but polar or charged solutes (such as NaCl) compensate for lost water-water hydrogen bonds by forming new solute-water interactions. The net change in enthalpy (ΔH) for dissolving these solutes is generally small.

In addition to requiring this input of energy, dissolving hydrophobic compounds in water produces a measurable decrease in entropy. Water molecules in the immediate vicinity of a nonpolar solute are constrained in their possible orientations, as they form a highly ordered cagelike shell around each solute molecule to maximize solvent-solvent hydrogen bonding. These water molecules are not as highly oriented as those in clathrates, crystalline compounds of nonpolar solutes and water, but the effect is the same in both cases: the ordering of water molecules reduces entropy. The number of ordered water molecules, and therefore the magnitude of the entropy decrease, is proportional to the surface area of the hydrophobic solute enclosed within the cage of water molecules. The free-energy change for dissolving a nonpolar solute in water is thus unfavorable: ΔG = ΔH − TΔS, where ΔH has a positive value, ΔS has a negative value, and ΔG is positive.

When an amphipathic compound (Table 2-1) is mixed with water, the polar, hydrophilic region interacts favorably with the water and tends to dissolve, but the nonpolar, hydrophobic region tends to avoid contact with the water (Fig. 2-7a). The nonpolar regions of the molecules cluster together to present the smallest hydrophobic area to the aqueous solvent, and the polar regions are arranged to maximize their interaction with each other and with the solvent (Fig. 2-7b), a phenomenon called the hydrophobic effect.

Many biomolecules are amphipathic; proteins, pigments, certain vitamins, and the sterols and phospholipids of membranes all have both polar and nonpolar surface regions. Structures composed of these molecules are stabilized by the hydrophobic effect, which favors aggregation of the nonpolar regions. The hydrophobic effect on interactions among lipids, and between lipids and proteins, is the most important determinant of structure in biological membranes. The aggregation of nonpolar amino acids in protein interiors, driven by the hydrophobic effect, also stabilizes the three-dimensional structures of proteins.

When two uncharged atoms are brought very close together, their surrounding electron clouds influence each other. Random variations in the positions of the electrons around one nucleus may create a transient electric dipole, which induces a transient, opposite electric dipole in the nearby atom. The two dipoles weakly attract each other, bringing the two nuclei closer. These weak attractions are called van der Waals interactions (also known as London dispersion forces). As the two nuclei draw closer together, their electron clouds begin to repel each other.

The noncovalent interactions we have described—hydrogen bonds and ionic, hydrophobic, and van der Waals interactions (Table 2-4)—are much weaker than covalent bonds. An input of about 350 kJ of energy is required to break a mole of (6 × 1023) C—C single bonds, and about 410 kJ is needed to break a mole of C—H bonds, but as little as 4 kJ is sufficient to disrupt a mole of typical van der Waals interactions. Interactions driven by the hydrophobic effect are also much weaker than covalent bonds, although they are substantially strengthened by a highly polar solvent (a concentrated salt solution, for example). Ionic interactions and hydrogen bonds are variable in strength, depending on the polarity of the solvent and the alignment of the hydrogen-bonded atoms, but they are always significantly weaker than covalent bonds.

Although these four types of interactions are individually weak relative to covalent bonds, the cumulative effect of many such interactions can be very significant. For example, the noncovalent binding of an enzyme to its substrate may involve several hydrogen bonds and one or more ionic interactions, as well as the hydrophobic effect and van der Waals interactions. The formation of each of these associations contributes to a net decrease in the free energy of the system. We can calculate the stability of a noncovalent interaction, such as the hydrogen bonding of a small molecule to its macromolecular partner, from the binding energy, the reduction in the energy of the system when binding occurs. Stability, as measured by the equilibrium constant (discussed in Section 2.2) of the binding reaction, varies exponentially with binding energy. To dissociate two biomolecules (such as an enzyme and its bound substrate) that are associated noncovalently through multiple weak interactions, all these interactions must be disrupted at the same time.

Macromolecules such as proteins, DNA, and RNA contain so many sites of potential hydrogen bonding or ionic, van der Waals, or hydrophobic clustering that the cumulative effect can be enormous. For macromolecules, the most stable (that is, the native) structure is usually that in which these weak interactions are maximized. The folding of a single polypeptide or polynucleotide chain into its three-dimensional shape is determined by this principle. The binding of an antigen to a specific antibody depends on the cumulative effects of many weak interactions. The energy released when an enzyme binds noncovalently to its substrate is the main source of the enzyme’s catalytic power. The binding of a hormone or a neurotransmitter to its cellular receptor protein is the result of multiple weak interactions.

The hemoglobin molecules carrying oxygen from your lungs to your brain at this moment were synthesized within the past month; by next month they will have been degraded and entirely replaced by new hemoglobin molecules. The glucose you ingested with your most recent meal is now circulating in your bloodstream; before the day is over these particular glucose molecules will have been converted into something else — carbon dioxide or fat, perhaps — and will have been replaced with a fresh supply of glucose, so that your blood glucose concentration is more or less constant over the whole day. The amounts of hemoglobin and glucose in the blood remain nearly constant because the rate of synthesis or intake of each just balances the rate of its breakdown, consumption, or conversion into some other product. The constancy of concentration is the result of a dynamic steady state, a steady state that is far from equilibrium.

For chemical reactions occurring in solution, we can define a system as all the constituent reactants and products, the solvent that contains them, and the immediate atmosphere — in short, everything within a defined region of space. The system and its surroundings together constitute the universe. If the system exchanges neither matter nor energy with its surroundings, it is said to be isolated. If the system exchanges energy but not matter with its surroundings, it is a closed system; if it exchanges both energy and matter with its surroundings, it is an open system. A living organism is an open system; it exchanges both matter and energy with its surroundings. Organisms obtain energy from their surroundings in two ways: (1) they take up chemical fuels (such as glucose) from the environment and extract energy by oxidizing them (see Box 1-3, Case 2); or (2) they absorb energy from sunlight.

The randomness or disorder of the components of a chemical system is expressed as entropy, S (Box 1-3). Any change in randomness of the system is expressed as entropy change, ΔS, which by convention has a positive value when randomness increases. J. Willard Gibbs, the scientist who developed the theory of energy changes during chemical reactions, showed that the free energy, G, of any closed system can be defined in terms of three quantities: enthalpy, H, or heat content, roughly reflecting the number and kinds of bonds; entropy, S; and the absolute temperature, T (in Kelvin). The definition of free energy is G = H − TS.

To carry out these thermodynamically unfavorable, energy-requiring (endergonic) reactions, cells couple them to other reactions that liberate free energy (exergonic reactions), so that the overall process is exergonic: the sum of the free-energy changes is negative. The exergonic reaction most commonly employed in this way involves adenosine triphosphate (ATP; Fig. 1-25) in which two phosphoanhydride bonds are capable of supplying the free energy to make a coupled endergonic reaction possible. In Section 13.3 we discuss in more detail this role of ATP.

All biological macromolecules are much less thermodynamically stable than their monomeric subunits, yet they are kinetically stable: their uncatalyzed breakdown occurs so slowly (over years rather than seconds) that, on a time scale that matters for the organism, these molecules are stable. Virtually every chemical reaction in a cell occurs at a significant rate only because of the presence of enzymes — biocatalysts that, like all other catalysts, greatly enhance the rate of specific chemical reactions without being consumed in the process. The path from reactant(s) to product(s) almost invariably involves an energy barrier, called the activation barrier (Fig. 1-27), that must be surmounted for any reaction to proceed.

Other pathways start with small precursor molecules and convert them to progressively larger and more complex molecules, including proteins and nucleic acids. Such synthetic pathways, which invariably require the input of energy, are collectively designated anabolism. The overall network of enzyme-catalyzed pathways, both catabolic and anabolic, constitutes cellular metabolism. ATP (as well as other energetically equivalent nucleoside triphosphates) is the connecting link between the catabolic and anabolic components of this network (shown schematically in Fig. 1-28). The pathways of enzyme-catalyzed reactions that act on the main constituents of cells — proteins, fats, sugars, and nucleic acids — are nearly identical in all living organisms.

The capacity of living cells to preserve their genetic material and to duplicate it for the next generation results from the structural complementarity between the two strands of the DNA molecule (Fig. 1-30). The basic unit of DNA is a linear polymer of four different monomeric subunits, deoxyribonucleotides, arranged in a precise linear sequence. It is this linear sequence that encodes the genetic information. Two of these polymeric strands are twisted about each other to form the DNA double helix, in which each deoxyribonucleotide in one strand pairs specifically with a complementary deoxyribonucleotide in the opposite strand.

A linear sequence of deoxyribonucleotides in DNA codes (through an intermediary, RNA) for the production of a protein with a corresponding linear sequence of amino acids (Fig. 1-31). The protein folds into a particular three-dimensional shape, determined by its amino acid sequence and stabilized primarily by noncovalent interactions. Although the final shape of the folded protein is dictated by its amino acid sequence, the folding of many proteins is aided by “molecular chaperones” (see Fig. 4-28). The precise three-dimensional structure, or native conformation, of the protein is crucial to its function.

Once in its native conformation, a protein may associate noncovalently with other macromolecules (other proteins, nucleic acids, or lipids) to form supramolecular complexes such as chromosomes, ribosomes, and membranes. The individual molecules of these complexes have specific, high-affinity binding sites for each other, and within the cell they spontaneously self-assemble into functional complexes. Although the amino acid sequences of proteins carry all necessary information for achieving the proteins’ native conformation, accurate folding and self-assembly also require the right cellular environment — pH, ionic strength, metal ion concentrations, and so forth. Thus, DNA sequence alone is not enough to form and maintain a fully functioning cell.

The remarkable similarity of metabolic pathways and gene sequences across the three domains of life argues strongly that all modern organisms are derived from a common evolutionary progenitor by a series of small changes (mutations), each of which conferred a selective advantage to some organism in some ecological niche. Despite the near-perfect fidelity of genetic replication, infrequent unrepaired mistakes in the DNA replication process lead to changes in the nucleotide sequence of DNA, producing a genetic mutation and changing the instructions for a cellular component. Incorrectly repaired damage to one of the DNA strands has the same effect.

Occasionally, however, a mutation better equips an organism or cell to survive in its environment (Fig. 1-32). The mutant enzyme might have acquired a slightly different specificity, for example, so that it is now able to use some compound that the cell was previously unable to metabolize. If a population of cells were to find itself in an environment where that compound was the only or the most abundant available source of fuel, the mutant cell would have a selective advantage over the other, unmutated (wild-type) cells in the population. The mutant cell and its progeny would survive and prosper in the new environment, whereas wild-type cells would starve and be eliminated. This is what Charles Darwin meant by natural selection — what is sometimes summarized as “survival of the fittest.”

This hypothesis was tested in a classic experiment on the abiotic (nonbiological) origin of organic biomolecules carried out in 1953 by biochemist Stanley Miller in the laboratory of the physical chemist Harold Urey. Miller subjected gaseous mixtures such as those presumed to exist on the prebiotic Earth, including NH3, CH4, H2O, and H2, to electrical sparks produced across a pair of electrodes (to simulate lightning) for periods of a week or more, then analyzed the contents of the closed reaction vessel (Fig. 1-33). The gas phase of the resulting mixture contained CO and CO2 as well as the starting materials. The water phase contained a variety of organic compounds, including some amino acids, hydroxy acids, aldehydes, and hydrogen cyanide (HCN).

In modern organisms, nucleic acids encode the genetic information that specifies the structure of enzymes, and enzymes catalyze the replication and repair of nucleic acids. The mutual dependence of these two classes of biomolecules brings up the perplexing question: which came first, DNA or protein? The answer may be that they appeared about the same time, and RNA preceded them both. The discovery that RNA molecules can act as catalysts in their own formation suggests that RNA or a similar molecule may have been the first gene and the first catalyst.

The division of function between DNA (genetic information storage) and protein (catalysis) was, according to the “RNA world” hypothesis, a later development. New variants of self-replicating RNA molecules developed that had the additional ability to catalyze the condensation of amino acids into peptides. Occasionally, the peptide(s) thus formed would reinforce the self-replicating ability of the RNA, and the pair — RNA molecule and helping peptide — could undergo further modifications in sequence, generating increasingly efficient self-replicating systems. The remarkable discovery that in the protein-synthesizing machinery of modern cells (ribosomes), RNA molecules, not proteins, catalyze the formation of peptide bonds is consistent with the RNA world hypothesis.

Earth was formed about 4.6 billion years ago, and the first evidence of life dates to more than 3.5 billion years ago (see the timeline in Figure 1-36). In 1996, scientists working in Greenland found chemical evidence of life (“fossil molecules”) from as far back as 3.85 billion years ago, forms of carbon embedded in rock that seem to have a distinctly biological origin. Somewhere on Earth during its first billion years the first simple organism arose, capable of replicating its own structure from a template (RNA?) that was the first genetic material.