Chapter 2:  Cell Chemistry and Biosynthesis

Cells Are Made From a Few Types of Atoms

Each atom has at its center a positively charged nucleus, which is surrounded at some distance by a cloud of negatively charged electrons, held in a series of orbitals by electrostatic attraction to the nucleus. The nucleus in turn consists of two kinds of subatomic particles: protons, which are positively charged, and neutrons, which are electrically neutral. The number of protons in the atomic nucleus gives the atomic number. An atom of hydrogen has a nucleus composed of a single proton; so hydrogen, with an atomic number of 1, is the lightest element. An atom of carbon has six protons in its nucleus and an atomic number of 6 (Figure 2-1). The electric charge carried by each proton is exactly equal and opposite to the charge carried by a single electron. Since an atom as a whole is electrically neutral, the number of negatively charged electrons surrounding the nucleus is equal to the number of positively charged protons that the nucleus contains; thus the number of electrons in an atom also equals the atomic number. It is these electrons that determine the chemical behavior of an atom, and all of the atoms of a given element have the same atomic number.

Neutrons are uncharged subatomic particles of essentially the same mass as protons. They contribute to the structural stability of the nucleus—if there are too many or too few, the nucleus may disintegrate by radioactive decay—but they do not alter the chemical properties of the atom. Thus an element can exist in several physically distinguishable but chemically identical forms, called isotopes, each isotope having a different number of neutrons but the same number of protons. Multiple isotopes of almost all the elements occur naturally, including some that are unstable. For example, while most carbon on Earth exists as the stable isotope carbon 12, with six protons and six neutrons, there are also small amounts of an unstable isotope, the radioactive carbon 14, whose atoms have six protons and eight neutrons. Carbon 14 undergoes radioactive decay at a slow but steady rate. This forms the basis for a technique known as carbon 14 dating, which is used in archaeology to determine the time of origin of organic materials.

The atomic weight of an atom, or the molecular weight of a molecule, is its mass relative to that of a hydrogen atom. This is essentially equal to the number of protons plus neutrons that the atom or molecule contains, since the electrons are much lighter and contribute almost nothing to the total. Thus the major isotope of carbon has an atomic weight of 12 and is symbolized as ¹²C, whereas the unstable isotope just discussed has an atomic weight of 14 and is written as ¹⁴C. The mass of an atom or a molecule is often specified in daltons, one dalton being an atomic mass unit approximately equal to the mass of a hydrogen atom.

Atoms are so small that it is hard to imagine their size. An individual carbon atom is roughly 0.2 nm in diameter, so that it would take about 5 million of them, laid out in a straight line, to span a millimeter. One proton or neutron weighs approximately 1/(6 × 10²³) gram, so one gram of hydrogen contains 6 × 10²³ atoms. This huge number (6 × 10²³, called Avogadro's number) is the key scale factor describing the relationship between everyday quantities and quantities measured in terms of individual atoms or molecules. If a substance has a molecular weight of X, 6 × 10²³ molecules of it will have a mass of X grams. This quantity is called one mole of the substance (Figure 2-2).

There are 92 naturally occurring elements, each differing from the others in the number of protons and electrons in its atoms. Living organisms, however, are made of only a small selection of these elements, four of which—carbon (C), hydrogen (H), nitrogen (N), and oxygen (O)—make up 96.5% of an organism's weight. This composition differs markedly from that of the nonliving inorganic environment (Figure 2-3) and is evidence of a distinctive type of chemistry. The most common elements in living organisms are listed in Table 2-1 with some of their atomic characteristics.

The Outermost Electrons Determine How Atoms Interact

To understand how atoms bond together to form the molecules that make up living organisms, we have to pay special attention to their electrons. Protons and neutrons are welded tightly to one another in the nucleus and change partners only under extreme conditions—during radioactive decay, for example, or in the interior of the sun or of a nuclear reactor. In living tissues, it is only the electrons of an atom that undergo rearrangements. They form the exterior of an atom and specify the rules of chemistry by which atoms combine to form molecules.

Electrons are in continuous motion around the nucleus, but motions on this submicroscopic scale obey different laws from those we are familiar with in everyday life. These laws dictate that electrons in an atom can exist only in certain discrete states, called orbitals, and that there is a strict limit to the number of electrons that can be accommodated in an orbital of a given type—a so-called electron shell. The electrons closest on average to the positive nucleus are attracted most strongly to it and occupy the innermost, most tightly bound shell. This shell can hold a maximum of two electrons. The second shell is farther away from the nucleus, and its electrons are less tightly bound. This second shell can hold up to eight electrons. The third shell contains electrons that are even less tightly bound; it can also hold up to eight electrons. The fourth and fifth shells can hold 18 electrons each. Atoms with more than four shells are very rare in biological molecules.

The electron arrangement of an atom is most stable when all the electrons are in the most tightly bound states that are possible for them—that is, when they occupy the innermost shells. Therefore, with certain exceptions in the larger atoms, the electrons of an atom fill the orbitals in order—the first shell before the second, the second before the third, and so on. An atom whose outermost shell is entirely filled with electrons is especially stable and therefore chemically unreactive. Examples are helium with 2 electrons, neon with 2 + 8, and argon with 2 + 8 + 8; these are all inert gases. Hydrogen, by contrast, with only one electron and therefore only a half-filled shell, is highly reactive. Likewise, the other atoms found in living tissues all have incomplete outer electron shells and are therefore able to donate, accept, or share electrons with each other to form both molecules and ions (Figure 2-4).

Because an unfilled electron shell is less stable than a filled one, atoms with incomplete outer shells have a strong tendency to interact with other atoms in a way that causes them to either gain or lose enough electrons to achieve a completed outermost shell. This electron exchange can be achieved either by transferring electrons from one atom to another or by sharing electrons between two atoms. These two strategies generate two types of chemical bonds between atoms: an ionic bond is formed when electrons are donated by one atom to another, whereas a covalent bond is formed when two atoms share a pair of electrons (Figure 2-5). Often, the pair of electrons is shared unequally, with a partial transfer between the atoms; this intermediate strategy results in a polar covalent bond, as we shall discuss later.

An H atom, which needs only one more electron to fill its shell, generally acquires it by electron sharing, forming one covalent bond with another atom; in many cases this bond is polar. The other most common elements in living cells—C, N, and O, with an incomplete second shell, and P and S, with an incomplete third shell (see Figure 2-4)—generally share electrons and achieve a filled outer shell of eight electrons by forming several covalent bonds. The number of electrons that an atom must acquire or lose (either by sharing or by transfer) to attain a filled outer shell is known as its valence.

The crucial role of the outer electron shell in determining the chemical properties of an element means that, when the elements are listed in order of their atomic number, there is a periodic recurrence of elements with similar properties: an element with, say, an incomplete second shell containing one electron will behave in much the same way as an element that has filled its second shell and has an incomplete third shell containing one electron. The metals, for example, have incomplete outer shells with just one or a few electrons, whereas, as we have just seen, the inert gases have full outer shells.

Ionic Bonds Form by the Gain and Loss of Electrons

Ionic bonds are most likely to be formed by atoms that have just one or two electrons in addition to a filled outer shell or are just one or two electrons short of acquiring a filled outer shell. They can often attain a completely filled outer electron shell more easily by transferring electrons to or from another atom than by sharing electrons. For example, from Figure 2-4 we see that a sodium (Na) atom, with atomic number 11, can strip itself down to a filled shell by giving up the single electron external to its second shell. By contrast, a chlorine (Cl) atom, with atomic number 17, can complete its outer shell by gaining just one electron. Consequently, if a Na atom encounters a Cl atom, an electron can jump from the Na to the Cl, leaving both atoms with filled outer shells. The offspring of this marriage between sodium, a soft and intensely reactive metal, and chlorine, a toxic green gas, is table salt (NaCl).

When an electron jumps from Na to Cl, both atoms become electrically charged ions. The Na atom that lost an electron now has one less electron than it has protons in its nucleus; it therefore has a single positive charge (Na⁺). The Cl atom that gained an electron now has one more electron than it has protons and has a single negative charge (Cl^-). Positive ions are called cations, and negative ions, anions. Ions can be further classified according to how many electrons are lost or gained. Thus sodium and potassium (K) have one electron to lose and form cations with a single positive charge (Na⁺ and K⁺), whereas magnesium and calcium have two electrons to lose and form cations with two positive charges (Mg²⁺ and Ca²⁺).

Because of their opposite charges, Na⁺ and Cl^- are attracted to each other and are thereby held together in an ionic bond. A salt crystal contains astronomical numbers of Na⁺ and Cl^- (about 2 × 10¹⁹ ions of each type in a crystal 1 mm across) packed together in a precise three-dimensional array with their opposite charges exactly balanced (Figure 2-6). Substances such as NaCl, which are held together solely by ionic bonds, are generally called salts rather than molecules. Ionic bonds are just one of several types of noncovalent bonds that can exist between atoms, and we shall meet other examples.

Because of a favorable interaction between water molecules and ions, ionic bonds are greatly weakened by water; thus many salts (including NaCl) are highly soluble in water—dissociating into individual ions (such as Na⁺ and Cl^-), each surrounded by a group of water molecules. In contrast, covalent bond strengths are not affected in this way.

Covalent Bonds Form by the Sharing of Electrons

All the characteristics of a cell depend on the molecules it contains. A molecule is defined as a cluster of atoms held together by covalent bonds; here electrons are shared between atoms to complete the outer shells, rather than being transferred between them. In the simplest possible molecule—a molecule of hydrogen (H₂)—two H atoms, each with a single electron, share two electrons, which is the number required to fill the first shell. These shared electrons form a cloud of negative charge that is densest between the two positively charged nuclei and helps to hold them together, in opposition to the mutual repulsion between like charges that would otherwise force them apart. The attractive and repulsive forces are in balance when the nuclei are separated by a characteristic distance, called the bond length.

A further crucial property of any bond—covalent or noncovalent—is its strength. Bond strength is measured by the amount of energy that must be supplied to break that bond. This is often expressed in units of kilocalories per mole (kcal/mole), where a kilocalorie is the amount of energy needed to raise the temperature of one liter of water by one degree centigrade. Thus if 1 kilocalorie must be supplied to break 6 × 10²³ bonds of a specific type (that is, 1 mole of these bonds), then the strength of that bond is 1 kcal/mole. An equivalent, widely used measure of energy is the kilojoule, which is equal to 0.239 kilocalories.

To get an idea of what bond strengths mean, it is helpful to compare them with the average energies of the impacts that molecules are constantly undergoing from collisions with other molecules in their environment (their thermal, or heat, energy), as well as with other sources of biological energy such as light and glucose oxidation (Figure 2-7). Typical covalent bonds are stronger than the thermal energies by a factor of 100, so they are resistant to being pulled apart by thermal motions and are normally broken only during specific chemical reactions with other atoms and molecules. The making and breaking of covalent bonds are violent events, and in living cells they are carefully controlled by highly specific catalysts, called enzymes. Noncovalent bonds as a rule are much weaker; we shall see later that they are important in the cell in the many situations where molecules have to associate and dissociate readily to carry out their functions.

Whereas an H atom can form only a single covalent bond, the other common atoms that form covalent bonds in cells—O, N, S, and P, as well as the all-important C atom—can form more than one. The outermost shell of these atoms, as we have seen, can accommodate up to eight electrons, and they form covalent bonds with as many other atoms as necessary to reach this number. Oxygen, with six electrons in its outer shell, is most stable when it acquires an extra two electrons by sharing with other atoms and therefore forms up to two covalent bonds. Nitrogen, with five outer electrons, forms a maximum of three covalent bonds, while carbon, with four outer electrons, forms up to four covalent bonds—thus sharing four pairs of electrons (see Figure 2-4).

When one atom forms covalent bonds with several others, these multiple bonds have definite orientations in space relative to one another, reflecting the orientations of the orbits of the shared electrons. Covalent bonds between multiple atoms are therefore characterized by specific bond angles as well as bond lengths and bond energies (Figure 2-8). The four covalent bonds that can form around a carbon atom, for example, are arranged as if pointing to the four corners of a regular tetrahedron. The precise orientation of covalent bonds forms the basis for the three-dimensional geometry of organic molecules.

There Are Different Types of Covalent Bonds

Most covalent bonds involve the sharing of two electrons, one donated by each participating atom; these are called single bonds. Some covalent bonds, however, involve the sharing of more than one pair of electrons. Four electrons can be shared, for example, two coming from each participating atom; such a bond is called a double bond. Double bonds are shorter and stronger than single bonds and have a characteristic effect on the three-dimensional geometry of molecules containing them. A single covalent bond between two atoms generally allows the rotation of one part of a molecule relative to the other around the bond axis. A double bond prevents such rotation, producing a more rigid and less flexible arrangement of atoms (Figure 2-9 and Panel 2-1, pp. 111–112).

Some molecules share electrons between three or more atoms, producing bonds that have a hybrid character intermediate between single and double bonds. The highly stable benzene molecule, for example, comprises a ring of six carbon atoms in which the bonding electrons are evenly distributed (although usually depicted as an alternating sequence of single and double bonds, as shown in Panel 2-1).

When the atoms joined by a single covalent bond belong to different elements, the two atoms usually attract the shared electrons to different degrees. Compared with a C atom, for example, O and N atoms attract electrons relatively strongly, whereas an H atom attracts electrons more weakly. By definition, a polar structure (in the electrical sense) is one with positive charge concentrated toward one end (the positive pole) and negative charge concentrated toward the other (the negative pole). Covalent bonds in which the electrons are shared unequally in this way are therefore known as polar covalent bonds (Figure 2-10). For example, the covalent bond between oxygen and hydrogen, -O-H, or between nitrogen and hydrogen, -N-H, is polar, whereas that between carbon and hydrogen, -C-H, has the electrons attracted much more equally by both atoms and is relatively nonpolar.

Polar covalent bonds are extremely important in biology because they create permanent dipoles that allow molecules to interact through electrical forces. Any large molecule with many polar groups will have a pattern of partial positive and negative charges on its surface. When such a molecule encounters a second molecule with a complementary set of charges, the two molecules will be attracted to each other by permanent dipole interactions that resemble (but are weaker than) the ionic bonds discussed previously for NaCl.

An Atom Often Behaves as if It Has a Fixed Radius

When a covalent bond forms between two atoms, the sharing of electrons brings the nuclei of these atoms unusually close together. But most of the atoms that are rapidly jostling each other in cells are located in separate molecules. What happens when two such atoms touch?

For simplicity and clarity, atoms and molecules are usually represented in a highly schematic way—either as a line drawing of the structural formula or as a ball and stick model. However, a more accurate representation can be obtained through the use of so-called space-filling models. Here a solid envelope is used to represent the radius of the electron cloud at which strong repulsive forces prevent a closer approach of any second, non-bonded atom—the so-called van der Waals radius for an atom. This is possible because the amount of repulsion increases very steeply as two such atoms approach each other closely. At slightly greater distances, any two atoms will experience a weak attractive force, known as a van der Waals attraction. As a result, there is a distance at which repulsive and attractive forces precisely balance to produce an energy minimum in each atom's interaction with an atom of a second, non-bonded element (Figure 2-11).

Depending on the intended purpose, we shall represent small molecules either as line drawings, ball and stick models, or space filling models throughout this book. For comparison, the water molecule is represented in all three ways in Figure 2-12. When dealing with very large molecules, such as proteins, we shall often need to further simplify the representation used (see, for example, Panel 3-2, pp. 138–139).

Water Is the Most Abundant Substance in Cells

Water accounts for about 70% of a cell's weight, and most intracellular reactions occur in an aqueous environment. Life on Earth began in the ocean, and the conditions in that primeval environment put a permanent stamp on the chemistry of living things. Life therefore hinges on the properties of water.

In each water molecule (H₂O) the two H atoms are linked to the O atom by covalent bonds (see Figure 2-12). The two bonds are highly polar because the O is strongly attractive for electrons, whereas the H is only weakly attractive. Consequently, there is an unequal distribution of electrons in a water molecule, with a preponderance of positive charge on the two H atoms and of negative charge on the O (see Figure 2-10). When a positively charged region of one water molecule (that is, one of its H atoms) comes close to a negatively charged region (that is, the O) of a second water molecule, the electrical attraction between them can result in a weak bond called a hydrogen bond. These bonds are much weaker than covalent bonds and are easily broken by the random thermal motions due to the heat energy of the molecules, so each bond lasts only an exceedingly short time. But the combined effect of many weak bonds is far from trivial. Each water molecule can form hydrogen bonds through its two H atoms to two other water molecules, producing a network in which hydrogen bonds are being continually broken and formed (Panel 2-2, pp. 112–113). It is only because of the hydrogen bonds that link water molecules together that water is a liquid at room temperature, with a high boiling point and high surface tension—rather than a gas.

Molecules, such as alcohols, that contain polar bonds and that can form hydrogen bonds with water dissolve readily in water. As mentioned previously, molecules carrying plus or minus charges (ions) likewise interact favorably with water. Such molecules are termed hydrophilic, meaning that they are water-loving. A large proportion of the molecules in the aqueous environment of a cell necessarily fall into this category, including sugars, DNA, RNA, and a majority of proteins. Hydrophobic (water-hating) molecules, by contrast, are uncharged and form few or no hydrogen bonds, and so do not dissolve in water. Hydrocarbons are an important example (see Panel 2-1, pp. 110–111). In these molecules the H atoms are covalently linked to C atoms by a largely nonpolar bond. Because the H atoms have almost no net positive charge, they cannot form effective hydrogen bonds to other molecules. This makes the hydrocarbon as a whole hydrophobic—a property that is exploited in cells, whose membranes are constructed from molecules that have long hydrocarbon tails, as we shall see in Chapter 10.

Some Polar Molecules Form Acids and Bases in Water

One of the simplest kinds of chemical reaction, and one that has profound significance in cells, takes place when a molecule possessing a highly polar covalent bond between a hydrogen and a second atom dissolves in water. The hydrogen atom in such a molecule has largely given up its electron to the companion atom and so exists as an almost naked positively charged hydrogen nucleus—in other words, a proton (H ⁺ ) . When the polar molecule becomes surrounded by water molecules, the proton is attracted to the partial negative charge on the O atom of an adjacent water molecule and can dissociate from its original partner to associate instead with the oxygen atoms of the water molecule to generate a hydronium ion (H ₃ O ⁺ ) (Figure 2-13A). The reverse reaction also takes place very readily, so one has to imagine an equilibrium state in which billions of protons are constantly flitting to and fro from one molecule in the solution to another.

Substances that release protons to form H₃O⁺ when they dissolve in water are termed acids. The higher the concentration of H₃O⁺, the more acidic the solution. H₃O⁺ is present even in pure water, at a concentration of 10^-7 M, as a result of the movement of protons from one water molecule to another (Figure 2-13B). By tradition, the H₃O⁺ concentration is usually referred to as the H⁺ concentration, even though most H⁺ in an aqueous solution is present as H₃O⁺. To avoid the use of unwieldy numbers, the concentration of H⁺ is expressed using a logarithmic scale called the pH scale, as illustrated in Panel 2-2 (pp. 112–113). Pure water has a pH of 7.0.

Because the proton of a hydronium ion can be passed readily to many types of molecules in cells, altering their character, the concentration of H₃O⁺ inside a cell (the acidity) must be closely regulated. Molecules that can give up protons will do so more readily if the concentration of H₃O⁺ in solution is low and will tend to receive them back if the concentration in solution is high.

The opposite of an acid is a base. Just as the defining property of an acid is that it donates protons to a water molecule so as to raise the concentration of H₃O⁺ ions, the defining property of a base is that it raises the concentration of hydroxyl (OH^-) ions—which are formed by removal of a proton from a water molecule. Thus sodium hydroxide (NaOH) is basic (the term alkaline is also used) because it dissociates in aqueous solution to form Na⁺ ions and OH^- ions. Another class of bases, especially important in living cells, are those that contain NH₂ groups. These groups can generate OH^- by taking a proton from water: -NH₂ + H₂O → -NH₃ ⁺ + OH^-.

Because an OH^- ion combines with a H₃O⁺ ion to form two water molecules, an increase in the OH^- concentration forces a decrease in the concentration of H₃O⁺, and vice versa. A pure solution of water contains an equally low concentration (10^-7 M) of both ions; it is neither acidic nor basic and is therefore said to be neutral with a pH of 7.0. The inside of cells is kept close to neutrality.

Four Types of Noncovalent Interactions Help Bring Molecules Together in Cells

In aqueous solutions, covalent bonds are 10 to 100 times stronger than the other attractive forces between atoms, allowing their connections to define the boundaries of one molecule from another. But much of biology depends on the specific binding of different molecules to each other. This binding is mediated by a group of noncovalent attractions that are individually quite weak, but whose bond energies can sum to create an effective force between two separate molecules. We have already introduced three of these noncovalent forces: ionic bonds, hydrogen bonds and van der Waals attractions. In Table 2-2, the strengths of these three types of bonds are compared to that of a typical covalent bond, both in the presence and the absence of water. Because of their fundamental importance in all biological systems, we shall summarize their properties here.

• Ionic bonds. These are purely electrostatic attractions between oppositely charged atoms. As we saw for NaCl, these forces are quite strong in the absence of water. However, the polar water molecules cluster around both fully charged ions and polar molecules that contain permanent dipoles (Figure 2-14). This greatly reduces the potential attractiveness of these charged species for each other (see Table 2-2).

• Hydrogen bonds. The structure of a typical hydrogen bond is illustrated in Figure 2-15. This bond represents a special form of polar interaction in which an electropositive hydrogen atom is partially shared by two electronegative atoms. Its hydrogen can be viewed as a proton that has partially dissociated from a donor atom, allowing it to be shared by a second acceptor atom. Unlike a typical electrostatic interaction, this bond is highly directional—being strongest when a straight line can be drawn between all three of the involved atoms. As already discussed, water weakens these bonds by forming competing hydrogen-bond interactions with the involved molecules (see Table 2-2).

• van der Waals attractions. The electron cloud around any nonpolar atom will fluctuate, producing a flickering dipole. Such dipoles will transiently induce an oppositely polarized flickering dipole in a nearby atom. This interaction generates an attraction between atoms that is very weak. But since many atoms can be simultaneously in contact when two surfaces fit closely, the net result is often significant. These so-called van der Waals attractions are not weakened by water (see Table 2-2).

The fourth effect that can play an important part in bringing molecules together in water is a hydrophobic force. This force is caused by a pushing of nonpolar surfaces out of the hydrogen-bonded water network, where they would physically interfere with the highly favorable interactions between water molecules. Because bringing two nonpolar surfaces together reduces their contact with water, the force is a rather nonspecific one. Nevertheless, we shall see in Chapter 3 that hydrophobic forces are central to the proper folding of protein molecules.

Panel 2-3 provides an overview of the four types of interactions just described. And Figure 2-16 illustrates, in a schematic way, how many such interactions can sum to hold together the matching surfaces of two macromolecules, even though each interaction by itself would be much too weak to be effective.

A Cell Is Formed from Carbon Compounds

Having looked at the ways atoms combine into small molecules and how these molecules behave in an aqueous environment, we now examine the main classes of small molecules found in cells and their biological roles. We shall see that a few basic categories of molecules, formed from a handful of different elements, give rise to all the extraordinary richness of form and behavior shown by living things.

If we disregard water, nearly all the molecules in a cell are based on carbon. Carbon is outstanding among all the elements in its ability to form large molecules; silicon is a poor second. Because it is small and has four electrons and four vacancies in its outermost shell, a carbon atom can form four covalent bonds with other atoms. Most important, one carbon atom can join to other carbon atoms through highly stable covalent C-C bonds to form chains and rings and hence generate large and complex molecules with no obvious upper limit to their size (see Panel 2-1, pp. 110–111). The small and large carbon compounds made by cells are called organic molecules.

Certain combinations of atoms, such as the methyl (-CH₃), hydroxyl (-OH), carboxyl (-COOH), carbonyl (-C=O), phosphate (-PO₃ ^2-), and amino (-NH₂) groups, occur repeatedly in organic molecules. Each such chemical group has distinct chemical and physical properties that influence the behavior of the molecule in which the group occurs. The most common chemical groups and some of their properties are summarized in Panel 2-1, pp. 110–111.

Cells Contain Four Major Families of Small Organic Molecules

The small organic molecules of the cell are carbon-based compounds that have molecular weights in the range 100 to 1000 and contain up to 30 or so carbon atoms. They are usually found free in solution and have many different fates. Some are used as monomer subunits to construct the giant polymeric macromolecules—the proteins, nucleic acids, and large polysaccharides—of the cell. Others act as energy sources and are broken down and transformed into other small molecules in a maze of intracellular metabolic pathways. Many small molecules have more than one role in the cell—for example, acting both as a potential subunit for a macromolecule and as an energy source. Small organic molecules are much less abundant than the organic macromolecules, accounting for only about one-tenth of the total mass of organic matter in a cell (Table 2-3). As a rough guess, there may be a thousand different kinds of these small molecules in a typical cell.

All organic molecules are synthesized from and are broken down into the same set of simple compounds. Both their synthesis and their breakdown occur through sequences of chemical changes that are limited in scope and follow definite rules. As a consequence, the compounds in a cell are chemically related and most can be classified into a small number of distinct families. Broadly speaking, cells contain four major families of small organic molecules: the sugars, the fatty acids, the amino acids, and the nucleotides (Figure 2-17). Although many compounds present in cells do not fit into these categories, these four families of small organic molecules, together with the macromolecules made by linking them into long chains, account for a large fraction of cell mass (see Table 2-3).

Sugars Provide an Energy Source for Cells and Are the Subunits of Polysaccharides

The simplest sugars—the monosaccharides—are compounds with the general formula (CH₂O)_n, where n is usually 3, 4, 5, 6, 7, or 8. Sugars, and the molecules made from them, are also called carbohydrates because of this simple formula. Glucose, for example, has the formula C₆H₁₂O₆ (Figure 2-18). The formula, however, does not fully define the molecule: the same set of carbons, hydrogens, and oxygens can be joined together by covalent bonds in a variety of ways, creating structures with different shapes. As shown in Panel 2-4 (pp. 116–117), for example, glucose can be converted into a different sugar—mannose or galactose—simply by switching the orientations of specific OH groups relative to the rest of the molecule. Each of these sugars, moreover, can exist in either of two forms, called the d-form and the l-form, which are mirror images of each other. Sets of molecules with the same chemical formula but different structures are called isomers, and the subset of such molecules that are mirror-image pairs are called optical isomers. Isomers are widespread among organic molecules in general, and they play a major part in generating the enormous variety of sugars.

An outline of sugar structures and chemistry is given in Panel 2-4. Sugars can exist in either a ring or an open-chain form. In their open-chain form, sugars contain a number of hydroxyl groups and either one aldehyde (_H>C=O) or one ketone ( C=O) group. The aldehyde or ketone group plays a special role. First, it can react with a hydroxyl group in the same molecule to convert the molecule into a ring; in the ring form the carbon of the original aldehyde or ketone group can be recognized as the only one that is bonded to two oxygens. Second, once the ring is formed, this same carbon can become further linked to one of the carbons bearing a hydroxyl group on another sugar molecule, creating a disaccharide; such as sucrose, which is composed of a glucose and a fructose unit. Larger sugar polymers range from the oligosaccharides (trisaccharides, tetrasaccharides, and so on) up to giant polysaccharides, which can contain thousands of monosaccharide units.

The way that sugars are linked together to form polymers illustrates some common features of biochemical bond formation. A bond is formed between an -OH group on one sugar and an -OH group on another by a condensation reaction, in which a molecule of water is expelled as the bond is formed (Figure 2-19). Subunits in other biological polymers, such as nucleic acids and proteins, are also linked by condensation reactions in which water is expelled. The bonds created by all of these condensation reactions can be broken by the reverse process of hydrolysis, in which a molecule of water is consumed (see Figure 2-19).

Because each monosaccharide has several free hydroxyl groups that can form a link to another monosaccharide (or to some other compound), sugar polymers can be branched, and the number of possible polysaccharide structures is extremely large. Even a simple disaccharide consisting of two glucose residues can exist in eleven different varieties (Figure 2-20), while three different hexoses (C₆H₁₂O₆) can join together to make several thousand trisaccharides. For this reason it is a much more complex task to determine the arrangement of sugars in a polysaccharide than to determine the nucleotide sequence of a DNA molecule, where each unit is joined to the next in exactly the same way.

The monosaccharide glucose has a central role as an energy source for cells. In a series of reactions, it is broken down to smaller molecules, releasing energy that the cell can harness to do useful work, as we shall explain later. Cells use simple polysaccharides composed only of glucose units—principally glycogen in animals and starch in plants—as long-term stores of energy.

Sugars do not function only in the production and storage of energy. They also can be used, for example, to make mechanical supports. Thus, the most abundant organic chemical on Earth—the cellulose of plant cell walls—is a polysaccharide of glucose. Another extraordinarily abundant organic substance, the chitin of insect exoskeletons and fungal cell walls, is also a polysaccharide—in this case a linear polymer of a sugar derivative called N-acetylglucosamine. Polysaccharides of various other sorts are the main components of slime, mucus, and gristle.

Smaller oligosaccharides can be covalently linked to proteins to form glycoproteins and to lipids to form glycolipids, which are found in cell membranes. As described in Chapter 10, the surfaces of most cells are clothed and decorated with sugar polymers belonging to glycoproteins and glycolipids in the cell membrane. These sugar side chains are often recognized selectively by other cells. And differences between people in the details of their cell-surface sugars are the molecular basis for the major different human blood groups.

Fatty Acids Are Components of Cell Membranes

A fatty acid molecule, such as palmitic acid, has two chemically distinct regions (Figure 2-21). One is a long hydrocarbon chain, which is hydrophobic and not very reactive chemically. The other is a carboxyl (-COOH) group, which behaves as an acid (carboxylic acid): it is ionized in solution (-COO^-), extremely hydrophilic, and chemically reactive. Almost all the fatty acid molecules in a cell are covalently linked to other molecules by their carboxylic acid group.

The hydrocarbon tail of palmitic acid is saturated: it has no double bonds between carbon atoms and contains the maximum possible number of hydrogens. Stearic acid, another one of the common fatty acids in animal fat, is also saturated. Some other fatty acids, such as oleic acid, have unsaturated tails, with one or more double bonds along their length. The double bonds create kinks in the molecules, interfering with their ability to pack together in a solid mass. It is this that accounts for the difference between hard (saturated) and soft (polyunsaturated) margarine. The many different fatty acids found in cells differ only in the length of their hydrocarbon chains and the number and position of the carbon-carbon double bonds (see Panel 2-5, pp. 118–119).

Fatty acids serve as a concentrated food reserve in cells, because they can be broken down to produce about six times as much usable energy, weight for weight, as glucose. They are stored in the cytoplasm of many cells in the form of droplets of triacylglycerol molecules, which consist of three fatty acid chains joined to a glycerol molecule (see Panel 2-5); these molecules are the animal fats found in meat, butter, and cream, and the plant oils like corn oil and olive oil. When required to provide energy, the fatty acid chains are released from triacylglycerols and broken down into two-carbon units. These two-carbon units are identical to those derived from the breakdown of glucose and they enter the same energy-yielding reaction pathways, as will be described later in this chapter.

Fatty acids and their derivatives such as triacylglycerols are examples of lipids. Lipids comprise a loosely defined collection of biological molecules with the common feature that they are insoluble in water, while being soluble in fat and organic solvents such as benzene. They typically contain either long hydrocarbon chains, as in the fatty acids and isoprenes, or multiple linked aromatic rings, as in the steroids.

The most important function of fatty acids in cells is in the construction of cell membranes. These thin sheets enclose all cells and surround their internal organelles. They are composed largely of phospholipids, which are small molecules that, like triacylglycerols, are constructed mainly from fatty acids and glycerol. In phospholipids the glycerol is joined to two fatty acid chains, however, rather than to three as in triacylglycerols. The “third” site on the glycerol is linked to a hydrophilic phosphate group, which is in turn attached to a small hydrophilic compound such as choline (see Panel 2-5). Each phospholipid molecule, therefore, has a hydrophobic tail composed of the two fatty acid chains and a hydrophilic head, where the phosphate is located. This gives them different physical and chemical properties from triacylglycerols, which are predominantly hydrophobic. Molecules like phospholipids, with both hydrophobic and hydrophilic regions, are termed amphipathic.

The membrane-forming property of phospholipids results from their amphipathic nature. Phospholipids will spread over the surface of water to form a monolayer of phospholipid molecules, with the hydrophobic tails facing the air and the hydrophilic heads in contact with the water. Two such molecular layers can readily combine tail-to-tail in water to make a phospholipid sandwich, or lipid bilayer. This bilayer is the structural basis of all cell membranes (Figure 2-22).

Amino Acids Are the Subunits of Proteins

Amino acids are a varied class of molecules with one defining property: they all possess a carboxylic acid group and an amino group, both linked to a single carbon atom called the α-carbon (Figure 2-23). Their chemical variety comes from the side chain that is also attached to the α-carbon. The importance of amino acids to the cell comes from their role in making proteins, which are polymers of amino acids joined head-to-tail in a long chain that is then folded into a three-dimensional structure unique to each type of protein. The covalent linkage between two adjacent amino acids in a protein chain is called a peptide bond; the chain of amino acids is also known as a polypeptide (Figure 2-24). Regardless of the specific amino acids from which it is made, the polypeptide has an amino (NH₂) group at one end (its N-terminus) and a carboxyl (COOH) group at its other end (its C-terminus). This gives it a definite directionality—a structural (as opposed to an electrical) polarity.

Twenty types of amino acids are found commonly in proteins, each with a different side chain attached to the α-carbon atom (see Panel 3-1, pp. 132–133). The same 20 amino acids occur over and over again in all proteins, whether from bacteria, plants, or animals. How this precise set of 20 amino acids came to be chosen is one of the mysteries surrounding the evolution of life; there is no obvious chemical reason why other amino acids could not have served just as well. But once the choice was established, it could not be changed; too much depended on it.

Like sugars, all amino acids, except glycine, exist as optical isomers in d- and l-forms (see Panel 3-1). But only l-forms are ever found in proteins (although d-amino acids occur as part of bacterial cell walls and in some antibiotics). The origin of this exclusive use of l-amino acids to make proteins is another evolutionary mystery.

The chemical versatility that the 20 standard amino acids provide is vitally important to the function of proteins. Five of the 20 amino acids have side chains that can form ions in solution and thereby can carry a charge (Figure 2-25). The others are uncharged; some are polar and hydrophilic, and some are nonpolar and hydrophobic. As we shall discuss in Chapter 3, the collective properties of the amino acid side chains underlie all the diverse and sophisticated functions of proteins.

Nucleotides Are the Subunits of DNA and RNA

A nucleotide is a molecule made up of a nitrogen-containing ring compound linked to a five-carbon sugar. This sugar can be either ribose or deoxyribose, and it carries one or more phosphate groups (Panel 2-6, pp. 120–121). Nucleotides containing ribose are known as ribonucleotides, and those containing deoxyribose as deoxyribonucleotides. The nitrogen-containing rings are generally referred to as bases for historical reasons: under acidic conditions they can each bind an H⁺ (proton) and thereby increase the concentration of OH^- ions in aqueous solution. There is a strong family resemblance between the different bases. Cytosine (C), thymine (T), and uracil (U) are called pyrimidines because they all derive from a six-membered pyrimidine ring; guanine (G) and adenine (A) are purine compounds, and they have a second, five-membered ring fused to the six-membered ring. Each nucleotide is named from the base it contains (see Panel 2-6).

Nucleotides can act as short-term carriers of chemical energy. Above all others, the ribonucleotide adenosine triphosphate, or ATP (Figure 2-26), is used to transfer energy in hundreds of different cellular reactions. ATP is formed through reactions that are driven by the energy released by the oxidative breakdown of foodstuffs. Its three phosphates are linked in series by two phosphoanhydride bonds, whose rupture releases large amounts of useful energy. The terminal phosphate group in particular is frequently split off by hydrolysis, often transferring a phosphate to other molecules and releasing energy that drives energy-requiring biosynthetic reactions (Figure 2-27). Other nucleotide derivatives serve as carriers for the transfer of other chemical groups, as will be described later.

The most fundamental role of nucleotides in the cell, however, is in the storage and retrieval of biological information. Nucleotides serve as building blocks for the construction of nucleic acids—long polymers in which nucleotide subunits are covalently linked by the formation of a phosphodiester bond between the phosphate group attached to the sugar of one nucleotide and a hydroxyl group on the sugar of the next nucleotide. Nucleic acid chains are synthesized from energy-rich nucleoside triphosphates by a condensation reaction that releases inorganic pyrophosphate during phosphodiester bond formation.

There are two main types of nucleic acids, differing in the type of sugar in their sugar-phosphate backbone. Those based on the sugar ribose are known as ribonucleic acids, or RNA, and contain the bases A, G, C, and U. Those based on deoxyribose (in which the hydroxyl at the 2′ position of the ribose carbon ring is replaced by a hydrogen (see Panel 2-6) are known as deoxyribonucleic acids, or DNA, and contain the bases A, G, C, and T (T is chemically similar to the U in RNA, merely adding the methyl group on the pyrimidine ring) (Figure 2-28). RNA usually occurs in cells in the form of a single polynucleotide chain, but DNA is virtually always in the form of a double-stranded molecule—a DNA double helix composed of two polynucleotide chains running antiparallel to each other and held together by hydrogen-bonding between the bases of the two chains.

The linear sequence of nucleotides in a DNA or an RNA encodes the genetic information of the cell. The ability of the bases in different nucleic acid molecules to recognize and pair with each other by hydrogen-bonding (called base-pairing)—G with C, and A with either T or U—underlies all of heredity and evolution, as explained in Chapter 4.

The Chemistry of Cells is Dominated by Macromolecules with Remarkable Properties

On a weight basis, macromolecules are by far the most abundant of the carbon-containing molecules in a living cell (Figure 2-29 and Table 2-4). They are the principal building blocks from which a cell is constructed and also the components that confer the most distinctive properties of living things. The macromolecules in cells are polymers that are constructed simply by covalently linking small organic molecules (called monomers, or subunits) into long chains (Figure 2-30). Yet they have many remarkable properties that could not have been predicted from their simple constituents.

Proteins are especially abundant and versatile. They perform thousands of distinct functions in cells. Many proteins serve as enzymes, the catalysts that direct the large number of covalent bond-making and bond-breaking reactions that the cell needs. All of the reactions whereby cells extract energy from food molecules are catalyzed by proteins serving as enzymes, for example, and an enzyme called ribulose bisphosphate carboxylase converts CO₂ to sugars in photosynthetic organisms, producing most of the organic matter needed for life on Earth. Other proteins are used to build structural components, such as tubulin, a protein that self-assembles to make the cell's long microtubules—or histones, proteins that compact the DNA in chromosomes. Yet other proteins act as molecular motors to produce force and movement, as in the case of myosin in muscle. Proteins can also have a wide variety of other functions, and we shall examine the molecular basis for many of them later in this book. Here we merely mention some general principles of macromolecular chemistry that make such functions possible.

Although the chemical reactions for adding subunits to each polymer are different in detail for proteins, nucleic acids, and polysaccharides, they share important features. Each polymer grows by the addition of a monomer onto the end of a growing polymer chain in a condensation reaction, in which a molecule of water is lost with each subunit added (see Figure 2-19). The stepwise polymerization of monomers into a long chain is a simple way to manufacture a large, complex molecule, since the subunits are added by the same reaction performed over and over again by the same set of enzymes. In a sense, the process resembles the repetitive operation of a machine in a factory—except in one crucial respect. Apart from some of the polysaccharides, most macromolecules are made from a set of monomers that are slightly different from one another—for example, the 20 different amino acids from which proteins are made. It is critical to life that the polymer chain is not assembled at random from these subunits; instead the subunits are added in a particular order, or sequence. The elaborate mechanisms that allow this to be accomplished by enzymes are described in detail in Chapters 5 and 6.

Noncovalent Bonds Specify Both the Precise Shape of a Macromolecule and its Binding to Other Molecules

Most of the covalent bonds in a macromolecule allow rotation of the atoms they join, so that the polymer chain has great flexibility. In principle, this allows a macromolecule to adopt an almost unlimited number of shapes, or conformations, as the polymer chain writhes and rotates under the influence of random thermal energy. However, the shapes of most biological macromolecules are highly constrained because of the many weak noncovalent bonds that form between different parts of the molecule. If these noncovalent bonds are formed in sufficient numbers, the polymer chain can strongly prefer one particular conformation, determined by the linear sequence of monomers in its chain. Virtually all protein molecules and many of the small RNA molecules found in cells fold tightly into one highly preferred conformation in this way (Figure 2-31).

The four types of noncovalent interactions important in biological molecules have been previously described in this chapter, and they are reviewed in Panel 2-3 (pp. 114–115). Although individually very weak, these interactions not only cooperate to fold biological macromolecules into unique shapes: they can also add up to create a strong attraction between two different molecules when these molecules fit together very closely, like a hand in a glove. This form of molecular interaction provides for great specificity, inasmuch as the multipoint contacts required for strong binding make it possible for a macromolecule to select out—through binding—just one of the many thousands of other types of molecules present inside a cell. Moreover, because the strength of the binding depends on the number of noncovalent bonds that are formed, interactions of almost any affinity are possible—allowing rapid dissociation when necessary.

Binding of this type underlies all biological catalysis, making it possible for proteins to function as enzymes. Noncovalent interactions also allow macromolecules to be used as building blocks for the formation of larger structures. In cells, macromolecules often bind together into large complexes, thereby forming intricate machines with multiple moving parts that perform such complex tasks as DNA replication and protein synthesis (Figure 2-32).

Summary

Living organisms are autonomous, self-propagating chemical systems. They are made from a distinctive and restricted set of small carbon-based molecules that are essentially the same for every living species. Each of these molecules is composed of a small set of atoms linked to each other in a precise configuration through covalent bonds. The main categories are sugars, fatty acids, amino acids, and nucleotides. Sugars are a primary source of chemical energy for cells and can be incorporated into polysaccharides for energy storage. Fatty acids are also important for energy storage, but their most critical function is in the formation of cell membranes. Polymers consisting of amino acids constitute the remarkably diverse and versatile macromolecules known as proteins. Nucleotides play a central part in energy transfer. They are also the subunits from which the informational macromolecules, RNA and DNA, are made.

Most of the dry mass of a cell consists of macromolecules that have been produced as linear polymers of amino acids (proteins) or nucleotides (DNA and RNA), covalently linked to each other in an exact order. The protein molecules and many of the RNAs fold into a unique conformation that depends on their sequence of subunits. This folding process creates unique surfaces, and it depends on a large set of weak interactions produced by noncovalent forces between atoms. These forces are of four types: ionic bonds, hydrogen bonds, van der Waals attractions, and an interaction between nonpolar groups caused by their hydrophobic expulsion from water. The same set of weak forces governs the specific binding of other molecules to macromolecules, making possible the myriad associations between biological molecules that produce the structure and the chemistry of a cell.

Cell Metabolism Is Organized by Enzymes

The chemical reactions that a cell carries out would normally occur only at temperatures that are much higher than those existing inside cells. For this reason, each reaction requires a specific boost in chemical reactivity. This requirement is crucial, because it allows each reaction to be controlled by the cell. The control is exerted through the specialized proteins called enzymes, each of which accelerates, or catalyzes, just one of the many possible kinds of reactions that a particular molecule might undergo. Enzyme-catalyzed reactions are usually connected in series, so that the product of one reaction becomes the starting material, or substrate, for the next (Figure 2-34). These long linear reaction pathways are in turn linked to one another, forming a maze of interconnected reactions that enable the cell to survive, grow, and reproduce (Figure 2-35).

Two opposing streams of chemical reactions occur in cells: (1) the catabolic pathways break down foodstuffs into smaller molecules, thereby generating both a useful form of energy for the cell and some of the small molecules that the cell needs as building blocks, and (2) the anabolic, or biosynthetic, pathways use the energy harnessed by catabolism to drive the synthesis of the many other molecules that form the cell. Together these two sets of reactions constitute the metabolism of the cell (Figure 2-36).

Many of the details of cell metabolism form the traditional subject of biochemistry and need not concern us here. But the general principles by which cells obtain energy from their environment and use it to create order are central to cell biology. We begin with a discussion of why a constant input of energy is needed to sustain living organisms.

Biological Order Is Made Possible by the Release of Heat Energy from Cells

The universal tendency of things to become disordered is expressed in a fundamental law of physics—the second law of thermodynamics—which states that in the universe, or in any isolated system (a collection of matter that is completely isolated from the rest of the universe), the degree of disorder can only increase. This law has such profound implications for all living things that it is worth restating in several ways.

For example, we can present the second law in terms of probability and state that systems will change spontaneously toward those arrangements that have the greatest probability. If we consider, for example, a box of 100 coins all lying heads up, a series of accidents that disturbs the box will tend to move the arrangement toward a mixture of 50 heads and 50 tails. The reason is simple: there is a huge number of possible arrangements of the individual coins in the mixture that can achieve the 50-50 result, but only one possible arrangement that keeps all of the coins oriented heads up. Because the 50-50 mixture is therefore the most probable, we say that it is more “disordered.” For the same reason, it is a common experience that one's living space will become increasingly disordered without intentional effort: the movement toward disorder is a spontaneous process, requiring a periodic effort to reverse it (Figure 2-37).

The amount of disorder in a system can be quantified. The quantity that we use to measure this disorder is called the entropy of the system: the greater the disorder, the greater the entropy. Thus, a third way to express the second law of thermodynamics is to say that systems will change spontaneously toward arrangements with greater entropy.

Living cells—by surviving, growing, and forming complex organisms—are generating order and thus might appear to defy the second law of thermodynamics. How is this possible? The answer is that a cell is not an isolated system: it takes in energy from its environment in the form of food, or as photons from the sun (or even, as in some chemosynthetic bacteria, from inorganic molecules alone), and it then uses this energy to generate order within itself. In the course of the chemical reactions that generate order, part of the energy that the cell uses is converted into heat. The heat is discharged into the cell's environment and disorders it, so that the total entropy—that of the cell plus its surroundings—increases, as demanded by the laws of physics.

To understand the principles governing these energy conversions, think of a cell as sitting in a sea of matter representing the rest of the universe. As the cell lives and grows, it creates internal order. But it releases heat energy as it synthesizes molecules and assembles them into cell structures. Heat is energy in its most disordered form—the random jostling of molecules. When the cell releases heat to the sea, it increases the intensity of molecular motions there (thermal motion)—thereby increasing the randomness, or disorder, of the sea. The second law of thermodynamics is satisfied because the increase in the amount of order inside the cell is more than compensated by a greater decrease in order (increase in entropy) in the surrounding sea of matter (Figure 2-38).

Where does the heat that the cell releases come from? Here we encounter another important law of thermodynamics. The first law of thermodynamics states that energy can be converted from one form to another, but that it cannot be created or destroyed. Some forms of energy are illustrated in Figure 2-39. The amount of energy in different forms will change as a result of the chemical reactions inside the cell, but the first law tells us that the total amount of energy must always be the same. For example, an animal cell takes in foodstuffs and converts some of the energy present in the chemical bonds between the atoms of these food molecules (chemical bond energy) into the random thermal motion of molecules (heat energy). This conversion of chemical energy into heat energy is essential if the reactions inside the cell are to cause the universe as a whole to become more disordered—as required by the second law.

The cell cannot derive any benefit from the heat energy it releases unless the heat-generating reactions inside the cell are directly linked to the processes that generate molecular order. It is the tight coupling of heat production to an increase in order that distinguishes the metabolism of a cell from the wasteful burning of fuel in a fire. Later in this chapter, we shall illustrate how this coupling occurs. For the moment, it is sufficient to recognize that a direct linkage of the “burning” of food molecules to the generation of biological order is required if cells are to be able to create and maintain an island of order in a universe tending toward chaos.

Photosynthetic Organisms Use Sunlight to Synthesize Organic Molecules

All animals live on energy stored in the chemical bonds of organic molecules made by other organisms, which they take in as food. The molecules in food also provide the atoms that animals need to construct new living matter. Some animals obtain their food by eating other animals. But at the bottom of the animal food chain are animals that eat plants. The plants, in turn, trap energy directly from sunlight. As a result, all of the energy used by animal cells is derived ultimately from the sun.

Solar energy enters the living world through photosynthesis in plants and photosynthetic bacteria. Photosynthesis allows the electromagnetic energy in sunlight to be converted into chemical bond energy in the cell. Plants are able to obtain all the atoms they need from inorganic sources: carbon from atmospheric carbon dioxide, hydrogen and oxygen from water, nitrogen from ammonia and nitrates in the soil, and other elements needed in smaller amounts from inorganic salts in the soil. They use the energy they derive from sunlight to build these atoms into sugars, amino acids, nucleotides, and fatty acids. These small molecules in turn are converted into the proteins, nucleic acids, polysaccharides, and lipids that form the plant. All of these substances serve as food molecules for animals, if the plants are later eaten.

The reactions of photosynthesis take place in two stages (Figure 2-40). In the first stage, energy from sunlight is captured and transiently stored as chemical bond energy in specialized small molecules that act as carriers of energy and reactive chemical groups. (We discuss these activated carrier molecules later.) Molecular oxygen (O₂ gas) derived from the splitting of water by light is released as a waste product of this first stage.

In the second stage, the molecules that serve as energy carriers are used to help drive a carbon fixation process in which sugars are manufactured from carbon dioxide gas (CO₂) and water (H₂O), thereby providing a useful source of stored chemical bond energy and materials—both for the plant itself and for any animals that eat it. We describe the elegant mechanisms that underlie these two stages of photosynthesis in Chapter 14.

The net result of the entire process of photosynthesis, so far as the green plant is concerned, can be summarized simply in the equation

The sugars produced are then used both as a source of chemical bond energy and as a source of materials to make the many other small and large organic molecules that are essential to the plant cell.

Cells Obtain Energy by the Oxidation of Organic Molecules

All animal and plant cells are powered by energy stored in the chemical bonds of organic molecules, whether these be sugars that a plant has photosynthesized as food for itself or the mixture of large and small molecules that an animal has eaten. In order to use this energy to live, grow, and reproduce, organisms must extract it in a usable form. In both plants and animals, energy is extracted from food molecules by a process of gradual oxidation, or controlled burning.

The Earth's atmosphere contains a great deal of oxygen, and in the presence of oxygen the most energetically stable form of carbon is as CO₂ and that of hydrogen is as H₂O. A cell is therefore able to obtain energy from sugars or other organic molecules by allowing their carbon and hydrogen atoms to combine with oxygen to produce CO₂ and H₂O, respectively—a process called respiration.

Photosynthesis and respiration are complementary processes (Figure 2-41). This means that the transactions between plants and animals are not all one way. Plants, animals, and microorganisms have existed together on this planet for so long that many of them have become an essential part of the others' environments. The oxygen released by photosynthesis is consumed in the combustion of organic molecules by nearly all organisms. And some of the CO₂ molecules that are fixed today into organic molecules by photosynthesis in a green leaf were yesterday released into the atmosphere by the respiration of an animal—or by that of a fungus or bacterium decomposing dead organic matter. We therefore see that carbon utilization forms a huge cycle that involves the biosphere (all of the living organisms on Earth) as a whole, crossing boundaries between individual organisms (Figure 2-42). Similarly, atoms of nitrogen, phosphorus, and sulfur move between the living and nonliving worlds in cycles that involve plants, animals, fungi, and bacteria.

Oxidation and Reduction Involve Electron Transfers

The cell does not oxidize organic molecules in one step, as occurs when organic material is burned in a fire. Through the use of enzyme catalysts, metabolism takes the molecules through a large number of reactions that only rarely involve the direct addition of oxygen. Before we consider some of these reactions and the purpose behind them, we need to discuss what is meant by the process of oxidation.

Oxidation, in the sense used above, does not mean only the addition of oxygen atoms; rather, it applies more generally to any reaction in which electrons are transferred from one atom to another. Oxidation in this sense refers to the removal of electrons, and reduction—the converse of oxidation—means the addition of electrons. Thus, Fe²⁺ is oxidized if it loses an electron to become Fe³⁺, and a chlorine atom is reduced if it gains an electron to become Cl^-. Since the number of electrons is conserved (no loss or gain) in a chemical reaction, oxidation and reduction always occur simultaneously: that is, if one molecule gains an electron in a reaction (reduction), a second molecule loses the electron (oxidation). When a sugar molecule is oxidized to CO₂ and H₂O, for example, the O₂ molecules involved in forming H₂O gain electrons and thus are said to have been reduced.

The terms “oxidation” and “reduction” apply even when there is only a partial shift of electrons between atoms linked by a covalent bond (Figure 2-43). When a carbon atom becomes covalently bonded to an atom with a strong affinity for electrons, such as oxygen, chlorine, or sulfur, for example, it gives up more than its equal share of electrons and forms a polar covalent bond: the positive charge of the carbon nucleus is now somewhat greater than the negative charge of its electrons, and the atom therefore acquires a partial positive charge and is said to be oxidized. Conversely, a carbon atom in a C-H linkage has slightly more than its share of electrons, and so it is said to be reduced (see Figure 2-43).

When a molecule in a cell picks up an electron (e^-), it often picks up a proton (H⁺) at the same time (protons being freely available in water). The net effect in this case is to add a hydrogen atom to the molecule

Even though a proton plus an electron is involved (instead of just an electron), such hydrogenation reactions are reductions, and the reverse, dehydrogenation reactions, are oxidations. It is especially easy to tell whether an organic molecule is being oxidized or reduced: reduction is occurring if its number of C-H bonds increases, whereas oxidation is occurring if its number of C-H bonds decreases (see Figure 2-43B).

Cells use enzymes to catalyze the oxidation of organic molecules in small steps, through a sequence of reactions that allows useful energy to be harvested. We now need to explain how enzymes work and some of the constraints under which they operate.

Enzymes Lower the Barriers That Block Chemical Reactions

Consider the reaction

The paper burns readily, releasing to the atmosphere both energy as heat and water and carbon dioxide as gases, but the smoke and ashes never spontaneously retrieve these entities from the heated atmosphere and reconstitute themselves into paper. When the paper burns, its chemical energy is dissipated as heat—not lost from the universe, since energy can never be created or destroyed, but irretrievably dispersed in the chaotic random thermal motions of molecules. At the same time, the atoms and molecules of the paper become dispersed and disordered. In the language of thermodynamics, there has been a loss of free energy, that is, of energy that can be harnessed to do work or drive chemical reactions. This loss reflects a loss of orderliness in the way the energy and molecules were stored in the paper. We shall discuss free energy in more detail shortly, but the general principle is clear enough intuitively: chemical reactions proceed only in the direction that leads to a loss of free energy; in other words, the spontaneous direction for any reaction is the direction that goes “downhill.” A “downhill” reaction in this sense is often said to be energetically favorable.

Although the most energetically favorable form of carbon under ordinary conditions is as CO₂, and that of hydrogen is as H₂O, a living organism does not disappear in a puff of smoke, and the book in your hands does not burst into flames. This is because the molecules both in the living organism and in the book are in a relatively stable state, and they cannot be changed to a state of lower energy without an input of energy: in other words, a molecule requires activation energy—a kick over an energy barrier—before it can undergo a chemical reaction that leaves it in a more stable state (Figure 2-44). In the case of a burning book, the activation energy is provided by the heat of a lighted match. For the molecules in the watery solution inside a cell, the kick is delivered by an unusually energetic random collision with surrounding molecules—collisions that become more violent as the temperature is raised.

In a living cell, the kick over the energy barrier is greatly aided by a specialized class of proteins—the enzymes. Each enzyme binds tightly to one or two molecules, called substrates, and holds them in a way that greatly reduces the activation energy of a particular chemical reaction that the bound substrates can undergo. A substance that can lower the activation energy of a reaction is termed a catalyst; catalysts increase the rate of chemical reactions because they allow a much larger proportion of the random collisions with surrounding molecules to kick the substrates over the energy barrier, as illustrated in Figure 2-45. Enzymes are among the most effective catalysts known, speeding up reactions by a factor of as much as 10¹⁴, and they thereby allow reactions that would not otherwise occur to proceed rapidly at normal temperatures.

Enzymes are also highly selective. Each enzyme usually catalyzes only one particular reaction: in other words, it selectively lowers the activation energy of only one of the several possible chemical reactions that its bound substrate molecules could undergo. In this way, enzymes direct each of the many different molecules in a cell along specific reaction pathways (Figure 2-46).

The success of living organisms is attributable to a cell's ability to make enzymes of many types, each with precisely specified properties. Each enzyme has a unique shape containing an active site, a pocket or groove in the enzyme into which only particular substrates will fit (Figure 2-47). Like all other catalysts, enzyme molecules themselves remain unchanged after participating in a reaction and therefore can function over and over again. In Chapter 3, we discuss further how enzymes work, after we have looked in detail at the molecular structure of proteins.

How Enzymes Find Their Substrates: The Importance of Rapid Diffusion

A typical enzyme will catalyze the reaction of about a thousand substrate molecules every second. This means that it must be able to bind a new substrate molecule in a fraction of a millisecond. But both enzymes and their substrates are present in relatively small numbers in a cell. How do they find each other so fast? Rapid binding is possible because the motions caused by heat energy are enormously fast at the molecular level. These molecular motions can be classified broadly into three kinds: (1) the movement of a molecule from one place to another (translational motion), (2) the rapid back-and-forth movement of covalently linked atoms with respect to one another (vibrations), and (3) rotations. All of these motions are important in bringing the surfaces of interacting molecules together.

These rates of molecular motions can be measured by a variety of spectroscopic techniques. These indicate that a large globular protein is constantly tumbling, rotating about its axis about a million times per second. Molecules are also in constant translational motion, which causes them to explore the space inside the cell very efficiently by wandering through it—a process called diffusion. In this way, every molecule in a cell collides with a huge number of other molecules each second. As the molecules in a liquid collide and bounce off one another, an individual molecule moves first one way and then another, its path constituting a random walk (Figure 2-48). In such a walk, the average distance that each molecule travels (as the crow flies) from its starting point is proportional to the square root of the time involved: that is, if it takes a molecule 1 second on average to travel 1 μm, it takes 4 seconds to travel 2 μm, 100 seconds to travel 10 μm, and so on.

The inside of a cell is very crowded (Figure 2-49). Nevertheless, experiments in which fluorescent dyes and other labeled molecules are injected into cells show that small organic molecules diffuse through the watery gel of the cytosol nearly as rapidly as they do through water. A small organic molecule, for example, takes only about one-fifth of a second on average to diffuse a distance of 10 μm. Diffusion is therefore an efficient way for small molecules to move the limited distances in the cell (a typical animal cell is 15 μm in diameter).

Since enzymes move more slowly than substrates in cells, we can think of them as sitting still. The rate of encounter of each enzyme molecule with its substrate will depend on the concentration of the substrate molecule. For example, some abundant substrates are present at a concentration of 0.5 mM. Since pure water is 55 M, there is only about one such substrate molecule in the cell for every 10⁵ water molecules. Nevertheless, the active site on an enzyme molecule that binds this substrate will be bombarded by about 500,000 random collisions with the substrate molecule per second. (For a substrate concentration tenfold lower, the number of collisions drops to 50,000 per second, and so on.) A random encounter between the surface of an enzyme and the matching surface of its substrate molecule often leads immediately to the formation of an enzyme-substrate complex that is ready to react. A reaction in which a covalent bond is broken or formed can now occur extremely rapidly. When one appreciates how quickly molecules move and react, the observed rates of enzymatic catalysis do not seem so amazing.

Once an enzyme and substrate have collided and snuggled together properly at the active site, they form multiple weak bonds with each other that persist until random thermal motion causes the molecules to dissociate again. In general, the stronger the binding of the enzyme and substrate, the slower their rate of dissociation. However, when two colliding molecules have poorly matching surfaces, few noncovalent bonds are formed and their total energy is negligible compared with that of thermal motion. In this case the two molecules dissociate as rapidly as they come together. This is what prevents incorrect and unwanted associations from forming between mismatched molecules, such as between an enzyme and the wrong substrate.

The Free-Energy Change for a Reaction Determines Whether It Can Occur

We must now digress briefly to introduce some fundamental chemistry. Cells are chemical systems that must obey all chemical and physical laws. Although enzymes speed up reactions, they cannot by themselves force energetically unfavorable reactions to occur. In terms of a water analogy, enzymes by themselves cannot make water run uphill. Cells, however, must do just that in order to grow and divide: they must build highly ordered and energy-rich molecules from small and simple ones. We shall see that this is done through enzymes that directly couple energetically favorable reactions, which release energy and produce heat, to energetically unfavorable reactions, which produce biological order.

Before examining how such coupling is achieved, we must consider more carefully the term “energetically favorable.” According to the second law of thermodynamics, a chemical reaction can proceed spontaneously only if it results in a net increase in the disorder of the universe (see Figure 2-38). The criterion for an increase in disorder of the universe can be expressed most conveniently in terms of a quantity called the free energy, G , of a system. The value of G is of interest only when a system undergoes a change, and the change in G, denoted Δ G (delta G), is critical. Suppose that the system being considered is a collection of molecules. As explained in Panel 2-7 (pp. 122–123), free energy has been defined such that ΔG directly measures the amount of disorder created in the universe when a reaction takes place that involves these molecules. Energetically favorable reactions, by definition, are those that decrease free energy, or, in other words, have a negative ΔG and disorder the universe (Figure 2-50).

A familiar example of an energetically favorable reaction on a macroscopic scale is the “reaction” by which a compressed spring relaxes to an expanded state, releasing its stored elastic energy as heat to its surroundings; an example on a microscopic scale is the dissolving of salt in water. Conversely, energetically unfavorable reactions, with a positive ΔG—such as those in which two amino acids are joined together to form a peptide bond—by themselves create order in the universe. Therefore, these reactions can take place only if they are coupled to a second reaction with a negative ΔG so large that the ΔG of the entire process is negative (Figure 2-51).

The Concentration of Reactants Influences ΔG

As we have just described, a reaction A B will go in the direction A → B when the associated free-energy change, ΔG, is negative, just as a tensed spring left to itself will relax and lose its stored energy to its surroundings as heat. For a chemical reaction, however, ΔG depends not only on the energy stored in each individual molecule, but also on the concentrations of the molecules in the reaction mixture. Remember that ΔG reflects the degree to which a reaction creates a more disordered—in other words, a more probable—state of the universe. Recalling our coin analogy, it is very likely that a coin will flip from a head to a tail orientation if a jiggling box contains 90 heads and 10 tails, but this is a less probable event if the box contains 10 heads and 90 tails. For exactly the same reason, for a reversible reaction A B, a large excess of A over B will tend to drive the reaction in the direction A → B; that is, there will be a tendency for there to be more molecules making the transition A → B than there are molecules making the transition B → A. Therefore, the ΔG becomes more negative for the transition A → B (and more positive for the transition B → A) as the ratio of A to B increases.

How much of a concentration difference is needed to compensate for a given decrease in chemical bond energy (and accompanying heat release)? The answer is not intuitively obvious, but it can be determined from a thermodynamic analysis that makes it possible to separate the concentration-dependent and the concentration-independent parts of the free-energy change. The ΔG for a given reaction can thereby be written as the sum of two parts: the first, called the standard free-energy change, Δ G°, depends on the intrinsic characters of the reacting molecules; the second depends on their concentrations. For the simple reaction A → B at 37°C,

where ΔG is in kilocalories per mole, [A] and [B] denote the concentrations of A and B, ln is the natural logarithm, and 0.616 is RT—the product of the gas constant, R, and the abolute temperature, T.

Note that ΔG equals the value of ΔG° when the molar concentrations of A and B are equal (ln 1 = 0). As expected, ΔG becomes more negative as the ratio of B to A decreases (the ln of a number < 1 is negative).

Chemical equilibrium is reached when the concentration effect just balances the push given to the reaction by ΔG°, so that there is no net change of free energy to drive the reaction in either direction (Figure 2-52). Here ΔG = 0, and so the concentrations of A and B are such that

which means that there is chemical equilibrium at 37°C when

Table 2-5 shows how the equilibrium ratio of A to B (expressed as an equilibrium constant, K ) depends on the value of ΔG°.

It is important to recognize that when an enzyme (or any catalyst) lowers the activation energy for the reaction A → B, it also lowers the activation energy for the reaction B → A by exactly the same amount (see Figure 2-44). The forward and backward reactions will therefore be accelerated by the same factor by an enzyme, and the equilibrium point for the reaction (and ΔG°) remains unchanged (Figure 2-53).

For Sequential Reactions, ΔG° Values Are Additive

The course of most reactions can be predicted quantitatively. A large body of thermodynamic data has been collected that makes it possible to calculate the standard change in free energy, ΔG°, for most of the important metabolic reactions of the cell. The overall free-energy change for a metabolic pathway is then simply the sum of the free-energy changes in each of its component steps. Consider, for example, two sequential reactions

where the ΔG° values are +5 and -13 kcal/mole, respectively. (Recall that a mole is 6 × 10²³ molecules of a substance.) If these two reactions occur sequentially, the ΔG° for the coupled reaction will be -8 kcal/mole. Thus, the unfavorable reaction X → Y, which will not occur spontaneously, can be driven by the favorable reaction Y → Z, provided that the second reaction follows the first.

Cells can therefore cause the energetically unfavorable transition, X → Y, to occur if an enzyme catalyzing the X → Y reaction is supplemented by a second enzyme that catalyzes the energetically favorable reaction, Y → Z. In effect, the reaction Y → Z will then act as a “siphon” to drive the conversion of all of molecule X to molecule Y, and thence to molecule Z (Figure 2-54). For example, several of the reactions in the long pathway that converts sugars into CO₂ and H₂O would be energetically unfavorable if considered on their own. But the pathway nevertheless proceeds rapidly to completion because the total ΔG° for the series of sequential reactions has a large negative value.

But forming a sequential pathway is not adequate for many purposes. Often the desired pathway is simply X → Y, without further conversion of Y to some other product. Fortunately, there are other more general ways of using enzymes to couple reactions together. How these work is the topic we discuss next.

Activated Carrier Molecules are Essential for Biosynthesis

The energy released by the oxidation of food molecules must be stored temporarily before it can be channeled into the construction of other small organic molecules and of the larger and more complex molecules needed by the cell. In most cases, the energy is stored as chemical bond energy in a small set of activated “carrier molecules,” which contain one or more energy-rich covalent bonds. These molecules diffuse rapidly throughout the cell and thereby carry their bond energy from sites of energy generation to the sites where energy is used for biosynthesis and other needed cell activities (Figure 2-55).

The activated carriers store energy in an easily exchangeable form, either as a readily transferable chemical group or as high-energy electrons, and they can serve a dual role as a source of both energy and chemical groups in biosynthetic reactions. For historical reasons, these molecules are also sometimes referred to as coenzymes. The most important of the activated carrier molecules are ATP and two molecules that are closely related to each other, NADH and NADPH—as we discuss in detail shortly. We shall see that cells use activated carrier molecules like money to pay for reactions that otherwise could not take place.

The Formation of an Activated Carrier Is Coupled to an Energetically Favorable Reaction

When a fuel molecule such as glucose is oxidized in a cell, enzyme-catalyzed reactions ensure that a large part of the free energy that is released by oxidation is captured in a chemically useful form, rather than being released wastefully as heat. This is achieved by means of a coupled reaction, in which an energetically favorable reaction is used to drive an energetically unfavorable one that produces an activated carrier molecule or some other useful energy store. Coupling mechanisms require enzymes and are fundamental to all the energy trans-actions of the cell.

The nature of a coupled reaction is illustrated by a mechanical analogy in Figure 2-56, in which an energetically favorable chemical reaction is represented by rocks falling from a cliff. The energy of falling rocks would normally be entirely wasted in the form of heat generated by friction when the rocks hit the ground (see the falling brick diagram in Figure 2-39). By careful design, however, part of this energy could be used instead to drive a paddle wheel that lifts a bucket of water (Figure 2-56B). Because the rocks can now reach the ground only after moving the paddle wheel, we say that the energetically favorable reaction of rock falling has been directly coupled to the energetically unfavorable reaction of lifting the bucket of water. Note that because part of the energy is used to do work in (B), the rocks hit the ground with less velocity than in (A), and correspondingly less energy is wasted as heat.

Exactly analogous processes occur in cells, where enzymes play the role of the paddle wheel in our analogy. By mechanisms that will be discussed later in this chapter, they couple an energetically favorable reaction, such as the oxidation of foodstuffs, to an energetically unfavorable reaction, such as the generation of an activated carrier molecule. As a result, the amount of heat released by the oxidation reaction is reduced by exactly the amount of energy that is stored in the energy-rich covalent bonds of the activated carrier molecule. The activated carrier molecule in turn picks up a packet of energy of a size sufficient to power a chemical reaction elsewhere in the cell.

ATP Is the Most Widely Used Activated Carrier Molecule

The most important and versatile of the activated carriers in cells is ATP (adenosine triphosphate). Just as the energy stored in the raised bucket of water in Figure 2-56B can be used to drive a wide variety of hydraulic machines, ATP serves as a convenient and versatile store, or currency, of energy to drive a variety of chemical reactions in cells. ATP is synthesized in an energetically unfavorable phosphorylation reaction in which a phosphate group is added to ADP (adenosine diphosphate). When required, ATP gives up its energy packet through its energetically favorable hydrolysis to ADP and inorganic phosphate (Figure 2-57). The regenerated ADP is then available to be used for another round of the phosphorylation reaction that forms ATP.

The energetically favorable reaction of ATP hydrolysis is coupled to many otherwise unfavorable reactions through which other molecules are synthesized. We shall encounter several of these reactions later in this chapter. Many of them involve the transfer of the terminal phosphate in ATP to another molecule, as illustrated by the phosphorylation reaction in Figure 2-58.

ATP is the most abundant active carrier in cells. As one example, it is used to supply energy for many of the pumps that transport substances into and out of the cell (discussed in Chapter 11). It also powers the molecular motors that enable muscle cells to contract and nerve cells to transport materials from one end of their long axons to another (discussed in Chapter 16).

Energy Stored in ATP Is Often Harnessed to Join Two Molecules Together

We have previously discussed one way in which an energetically favorable reaction can be coupled to an energetically unfavorable reaction, X → Y, so as to enable it to occur. In that scheme a second enzyme catalyzes the energetically favorable reaction Y → Z, pulling all of the X to Y in the process (see Figure 2-54). But when the required product is Y and not Z, this mechanism is not useful.

A frequent type of reaction that is needed for biosynthesis is one in which two molecules, A and B, are joined together to produce A-B in the energetically unfavorable condensation reaction

There is an indirect pathway that allows A-H and B-OH to form A-B, in which a coupling to ATP hydrolysis makes the reaction go. Here energy from ATP hydrolysis is first used to convert B-OH to a higher-energy intermediate compound, which then reacts directly with A-H to give A-B. The simplest possible mechanism involves the transfer of a phosphate from ATP to B-OH to make B-OPO₃, in which case the reaction pathway contains only two steps:

The condensation reaction, which by itself is energetically unfavorable, is forced to occur by being directly coupled to ATP hydrolysis in an enzyme-catalyzed reaction pathway (Figure 2-59A).

A biosynthetic reaction of exactly this type is employed to synthesize the amino acid glutamine, as illustrated in Figure 2-59B. We will see shortly that very similar (but more complex) mechanisms are also used to produce nearly all of the large molecules of the cell.

NADH and NADPH Are Important Electron Carriers

Other important activated carrier molecules participate in oxidation-reduction reactions and are commonly part of coupled reactions in cells. These activated carriers are specialized to carry high-energy electrons and hydrogen atoms. The most important of these electron carriers are NAD ⁺ (nicotinamide adenine dinucleotide) and the closely related molecule NADP ⁺ (nicotinamide adenine dinucleotide phosphate). Later, we examine some of the reactions in which they participate. NAD⁺ and NADP⁺ each pick up a “packet of energy” corresponding to two high-energy electrons plus a proton (H⁺)—being converted to NADH (reduced nicotinamide adenine dinucleotide) and NADPH (reduced nicotinamide adenine dinucleotide phosphate), respectively. These molecules can therefore also be regarded as carriers of hydride ions (the H⁺ plus two electrons, or H^-).

Like ATP, NADPH is an activated carrier that participates in many important biosynthetic reactions that would otherwise be energetically unfavorable. The NADPH is produced according to the general scheme shown in Figure 2-60A. During a special set of energy-yielding catabolic reactions, a hydrogen atom plus two electrons are removed from the substrate molecule and added to the nicotinamide ring of NADP⁺ to form NADPH. This is a typical oxidation-reduction reaction; the substrate is oxidized and NADP⁺ is reduced. The structures of NADP⁺ and NADPH are shown in Figure 2-60B.

The hydride ion carried by NADPH is given up readily in a subsequent oxidation-reduction reaction, because the ring can achieve a more stable arrangement of electrons without it. In this subsequent reaction, which regenerates NADP⁺, it is the NADPH that becomes oxidized and the substrate that becomes reduced. The NADPH is an effective donor of its hydride ion to other molecules for the same reason that ATP readily transfers a phosphate: in both cases the transfer is accompanied by a large negative free-energy change. One example of the use of NADPH in biosynthesis is shown in Figure 2-61.

The difference of a single phosphate group has no effect on the electron-transfer properties of NADPH compared with NADH, but it is crucial for their distinctive roles. The extra phosphate group on NADPH is far from the region involved in electron transfer (see Figure 2-60B) and is of no importance to the transfer reaction. It does, however, give a molecule of NADPH a slightly different shape from that of NADH, and so NADPH and NADH bind as substrates to different sets of enzymes. Thus the two types of carriers are used to transfer electrons (or hydride ions) between different sets of molecules.

Why should there be this division of labor? The answer lies in the need to regulate two sets of electron-transfer reactions independently. NADPH operates chiefly with enzymes that catalyze anabolic reactions, supplying the high-energy electrons needed to synthesize energy-rich biological molecules. NADH, by contrast, has a special role as an intermediate in the catabolic system of reactions that generate ATP through the oxidation of food molecules, as we will discuss shortly. The genesis of NADH from NAD⁺ and that of NADPH from NADP⁺ occur by different pathways and are independently regulated, so that the cell can independently adjust the supply of electrons for these two contrasting purposes. Inside the cell the ratio of NAD⁺ to NADH is kept high, whereas the ratio of NADP⁺ to NADPH is kept low. This provides plenty of NAD⁺ to act as an oxidizing agent and plenty of NADPH to act as a reducing agent—as required for their special roles in catabolism and anabolism, respectively.

There Are Many Other Activated Carrier Molecules in Cells

Other activated carriers also pick up and carry a chemical group in an easily transferred, high-energy linkage (Table 2-6). For example, coenzyme A carries an acetyl group in a readily transferable linkage, and in this activated form is known as acetyl CoA (acetyl coenzyme A). The structure of acetyl CoA is illustrated in Figure 2-62; it is used to add two carbon units in the biosynthesis of larger molecules.

In acetyl CoA and the other carrier molecules in Table 2-6, the transferable group makes up only a small part of the molecule. The rest consists of a large organic portion that serves as a convenient “handle,” facilitating the recognition of the carrier molecule by specific enzymes. As with acetyl CoA, this handle portion very often contains a nucleotide, a curious fact that may be a relic from an early stage of evolution. It is currently thought that the main catalysts for early life-forms—before DNA or proteins—were RNA molecules (or their close relatives), as described in Chapter 6. It is tempting to speculate that many of the carrier molecules that we find today originated in this earlier RNA world, where their nucleotide portions could have been useful for binding them to RNA enzymes.

Examples of the type of transfer reactions catalyzed by the activated carrier molecules ATP (transfer of phosphate) and NADPH (transfer of electrons and hydrogen) have been presented in Figures 2-58 and 2-61, respectively. The reactions of other activated carrier molecules involve the transfers of methyl, carboxyl, or glucose group, for the purpose of biosynthesis. The activated carriers required are usually generated in reactions that are coupled to ATP hydrolysis, as in the example in Figure 2-63. Therefore, the energy that enables their groups to be used for biosynthesis ultimately comes from the catabolic reactions that generate ATP. Similar processes occur in the synthesis of the very large molecules of the cell—the nucleic acids, proteins, and polysaccharides—that we discuss next.

The Synthesis of Biological Polymers Requires an Energy Input

As discussed previously, the macromolecules of the cell constitute the vast majority of its dry mass—that is, of the mass not due to water (see Figure 2-29). These molecules are made from subunits (or monomers) that are linked together in a condensation reaction, in which the constituents of a water molecule (OH plus H) are removed from the two reactants. Consequently, the reverse reaction—the breakdown of all three types of polymers—occurs by the enzyme-catalyzed addition of water (hydrolysis). This hydrolysis reaction is energetically favorable, whereas the biosynthetic reactions require an energy input and are more complex (Figure 2-64).

The nucleic acids (DNA and RNA), proteins, and polysaccharides are all polymers that are produced by the repeated addition of a subunit (also called a monomer) onto one end of a growing chain. The synthesis reactions for these three types of macromolecules are outlined in Figure 2-65. As indicated, the condensation step in each case depends on energy from nucleoside triphosphate hydrolysis. And yet, except for the nucleic acids, there are no phosphate groups left in the final product molecules. How are the reactions that release the energy of ATP hydrolysis coupled to polymer synthesis?

For each type of macromolecule, an enzyme-catalyzed pathway exists which resembles that discussed previously for the synthesis of the amino acid glutamine (see Figure 2-59). The principle is exactly the same, in that the OH group that will be removed in the condensation reaction is first activated by becoming involved in a high-energy linkage to a second molecule. However, the actual mechanisms used to link ATP hydrolysis to the synthesis of proteins and polysaccharides are more complex than that used for glutamine synthesis, since a series of high-energy intermediates is required to generate the final high-energy bond that is broken during the condensation step (discussed in Chapter 6 for protein synthesis).

There are limits to what each activated carrier can do in driving biosynthesis. The ΔG for the hydrolysis of ATP to ADP and inorganic phosphate (P_i) depends on the concentrations of all of the reactants, but under the usual conditions in a cell it is between -11 and -13 kcal/mole. In principle, this hydrolysis reaction can be used to drive an unfavorable reaction with a ΔG of, perhaps, +10 kcal/mole, provided that a suitable reaction path is available. For some biosynthetic reactions, however, even -13 kcal/mole may not be enough. In these cases the path of ATP hydrolysis can be altered so that it initially produces AMP and pyrophosphate (PP_i), which is itself then hydrolyzed in a subsequent step (Figure 2-66). The whole process makes available a total free-energy change of about -26 kcal/mole. An important biosynthetic reaction that is driven in this way is nucleic acid (polynucleotide) synthesis, as illustrated in Figure 2-67.

It is interesting to note that the polymerization reactions that produce macromolecules can be oriented in one of two ways, giving rise to either the head polymerization or the tail polymerization of monomers. In head polymerization the reactive bond required for the condensation reaction is carried on the end of the growing polymer, and it must therefore be regenerated each time that a monomer is added. In this case, each monomer brings with it the reactive bond that will be used in adding the next monomer in the series. In tail polymerization the reactive bond carried by each monomer is instead used immediately for its own addition (Figure 2-68).

We shall see in later chapters that both these types of polymerization are used. The synthesis of polynucleotides and some simple polysaccharides occurs by tail polymerization, for example, whereas the synthesis of proteins occurs by a head polymerization process.

Summary

Living cells are highly ordered and need to create order within themselves in order to survive and grow. This is thermodynamically possible only because of a continual input of energy, part of which must be released from the cells to their environment as heat. The energy comes ultimately from the electromagnetic radiation of the sun, which drives the formation of organic molecules in photosynthetic organisms such as green plants. Animals obtain their energy by eating these organic molecules and oxidizing them in a series of enzyme-catalyzed reactions that are coupled to the formation of ATP—a common currency of energy in all cells.

To make possible the continual generation of order in cells, the energetically favorable hydrolysis of ATP is coupled to energetically unfavorable reactions. In the biosynthesis of macromolecules, this is accomplished by the transfer of phosphate groups to form reactive phosphorylated intermediates. Because the energetically unfavorable reaction now becomes energetically favorable, ATP hydrolysis is said to drive the reaction. Polymeric molecules such as proteins, nucleic acids, and polysaccharides are assembled from small activated precursor molecules by repetitive condensation reactions that are driven in this way. Other reactive molecules, called either active carriers or coenzymes, transfer other chemical groups in the course of biosynthesis: NADPH transfers hydrogen as a proton plus two electrons (a hydride ion), for example, whereas acetyl CoA transfers an acetyl group.

Food Molecules Are Broken Down in Three Stages to Produce ATP

The proteins, lipids, and polysaccharides that make up most of the food we eat must be broken down into smaller molecules before our cells can use them—either as a source of energy or as building blocks for other molecules. The breakdown processes must act on food taken in from outside, but not on the macromolecules inside our own cells. Stage 1 in the enzymatic breakdown of food molecules is therefore digestion, which occurs either in our intestine outside cells, or in a specialized organelle within cells, the lysosome. (A membrane that surrounds the lysosome keeps its digestive enzymes separated from the cytosol, as described in Chapter 13.) In either case, the large polymeric molecules in food are broken down during digestion into their monomer subunits—proteins into amino acids, polysaccharides into sugars, and fats into fatty acids and glycerol—through the action of enzymes. After digestion, the small organic molecules derived from food enter the cytosol of the cell, where their gradual oxidation begins. As illustrated in Figure 2-70, oxidation occurs in two further stages of cellular catabolism: stage 2 starts in the cytosol and ends in the major energy-converting organelle, the mitochondrion; stage 3 is entirely confined to the mitochondrion.

In stage 2 a chain of reactions called glycolysis converts each molecule of glucose into two smaller molecules of pyruvate. Sugars other than glucose are similarly converted to pyruvate after their conversion to one of the sugar intermediates in this glycolytic pathway. During pyruvate formation, two types of activated carrier molecules are produced—ATP and NADH. The pyruvate then passes from the cytosol into mitochondria. There, each pyruvate molecule is converted into CO₂ plus a two-carbon acetyl group—which becomes attached to coenzyme A (CoA), forming acetyl CoA, another activated carrier molecule (see Figure 2-62). Large amounts of acetyl CoA are also produced by the stepwise breakdown and oxidation of fatty acids derived from fats, which are carried in the bloodstream, imported into cells as fatty acids, and then moved into mitochondria for acetyl CoA production.

Stage 3 of the oxidative breakdown of food molecules takes place entirely in mitochondria. The acetyl group in acetyl CoA is linked to coenzyme A through a high-energy linkage, and it is therefore easily transferable to other molecules. After its transfer to the four-carbon molecule oxaloacetate, the acetyl group enters a series of reactions called the citric acid cycle. As we discuss shortly, the acetyl group is oxidized to CO₂ in these reactions, and large amounts of the electron carrier NADH are generated. Finally, the high-energy electrons from NADH are passed along an electron-transport chain within the mitochondrial inner membrane, where the energy released by their transfer is used to drive a process that produces ATP and consumes molecular oxygen (O₂). It is in these final steps that most of the energy released by oxidation is harnessed to produce most of the cell's ATP.

Because the energy to drive ATP synthesis in mitochondria ultimately derives from the oxidative breakdown of food molecules, the phosphorylation of ADP to form ATP that is driven by electron transport in the mitochondrion is known as oxidative phosphorylation. The fascinating events that occur within the mitochondrial inner membrane during oxidative phosphorylation are the major focus of Chapter 14.

Through the production of ATP, the energy derived from the breakdown of sugars and fats is redistributed as packets of chemical energy in a form convenient for use elsewhere in the cell. Roughly 10⁹ molecules of ATP are in solution in a typical cell at any instant, and in many cells, all this ATP is turned over (that is, used up and replaced) every 1–2 minutes.

In all, nearly half of the energy that could in theory be derived from the oxidation of glucose or fatty acids to H₂O and CO₂ is captured and used to drive the energetically unfavorable reaction P_i + ADP → ATP. (By contrast, a typical combustion engine, such as a car engine, can convert no more than 20% of the available energy in its fuel into useful work.) The rest of the energy is released by the cell as heat, making our bodies warm.

Glycolysis Is a Central ATP-producing Pathway

The most important process in stage 2 of the breakdown of food molecules is the degradation of glucose in the sequence of reactions known as glycolysis—from the Greek glukus, “sweet,” and lusis, “rupture.” Glycolysis produces ATP without the involvement of molecular oxygen (O₂ gas). It occurs in the cytosol of most cells, including many anaerobic microorganisms (those that can live without utilizing molecular oxygen). Glycolysis probably evolved early in the history of life, before the activities of photosynthetic organisms introduced oxygen into the atmosphere. During glycolysis, a glucose molecule with six carbon atoms is converted into two molecules of pyruvate, each of which contains three carbon atoms. For each molecule of glucose, two molecules of ATP are hydrolyzed to provide energy to drive the early steps, but four molecules of ATP are produced in the later steps. At the end of glycolysis, there is consequently a net gain of two molecules of ATP for each glucose molecule broken down.

The glycolytic pathway is presented in outline in Figure 2-71, and in more detail in Panel 2-8 (pp. 124–125). Glycolysis involves a sequence of 10 separate reactions, each producing a different sugar intermediate and each catalyzed by a different enzyme. Like most enzymes, these enzymes all have names ending in ase—like isomerase and dehydrogenase—which indicate the type of reaction they catalyze.

Although no molecular oxygen is involved in glycolysis, oxidation occurs, in that electrons are removed by NAD⁺ (producing NADH) from some of the carbons derived from the glucose molecule. The stepwise nature of the process allows the energy of oxidation to be released in small packets, so that much of it can be stored in activated carrier molecules rather than all of it being released as heat (see Figure 2-69). Thus, some of the energy released by oxidation drives the direct synthesis of ATP molecules from ADP and P_i, and some remains with the electrons in the high-energy electron carrier NADH.

Two molecules of NADH are formed per molecule of glucose in the course of glycolysis. In aerobic organisms (those that require molecular oxygen to live), these NADH molecules donate their electrons to the electron-transport chain described in Chapter 14, and the NAD⁺ formed from the NADH is used again for glycolysis (see step 6 in Panel 2-8, pp. 124–125).

Fermentations Allow ATP to Be Produced in the Absence of Oxygen

For most animal and plant cells, glycolysis is only a prelude to the third and final stage of the breakdown of food molecules. In these cells, the pyruvate formed at the last step of stage 2 is rapidly transported into the mitochondria, where it is converted into CO₂ plus acetyl CoA, which is then completely oxidized to CO₂ and H₂O.

In contrast, for many anaerobic organisms—which do not utilize molecular oxygen and can grow and divide without it—glycolysis is the principal source of the cell's ATP. This is also true for certain animal tissues, such as skeletal muscle, that can continue to function when molecular oxygen is limiting. In these anaerobic conditions, the pyruvate and the NADH electrons stay in the cytosol. The pyruvate is converted into products excreted from the cell—for example, into ethanol and CO₂ in the yeasts used in brewing and breadmaking, or into lactate in muscle. In this process, the NADH gives up its electrons and is converted back into NAD⁺. This regeneration of NAD⁺ is required to maintain the reactions of glycolysis (Figure 2-72).

Anaerobic energy-yielding pathways like these are called fermentations. Studies of the commercially important fermentations carried out by yeasts inspired much of early biochemistry. Work in the nineteenth century led in 1896 to the then startling recognition that these processes could be studied outside living organisms, in cell extracts. This revolutionary discovery eventually made it possible to dissect out and study each of the individual reactions in the fermentation process. The piecing together of the complete glycolytic pathway in the 1930s was a major triumph of biochemistry, and it was quickly followed by the recognition of the central role of ATP in cellular processes. Thus, most of the fundamental concepts discussed in this chapter have been understood for more than 50 years.

Glycolysis Illustrates How Enzymes Couple Oxidation to Energy Storage

We have previously used a “paddle wheel” analogy to explain how cells harvest useful energy from the oxidation of organic molecules by using enzymes to couple an energetically unfavorable reaction to an energetically favorable one (see Figure 2-56). Enzymes play the part of the paddle wheel in our analogy, and we now return to a step in glycolysis that we have previously discussed, in order to illustrate exactly how coupled reactions occur.

Two central reactions in glycolysis (steps 6 and 7) convert the three-carbon sugar intermediate glyceraldehyde 3-phosphate (an aldehyde) into 3-phosphoglycerate (a carboxylic acid). This entails the oxidation of an aldehyde group to a carboxylic acid group, which occurs in two steps. The overall reaction releases enough free energy to convert a molecule of ADP to ATP and to transfer two electrons from the aldehyde to NAD⁺ to form NADH, while still releasing enough heat to the environment to make the overall reaction energetically favorable (ΔG° for the overall reaction is -3.0 kcal/mole).

The pathway by which this remarkable feat is accomplished is outlined in Figure 2-73. The chemical reactions are guided by two enzymes to which the sugar intermediates are tightly bound. The first enzyme (glyceraldehyde 3-phosphate dehydrogenase) forms a short-lived covalent bond to the aldehyde through a reactive -SH group on the enzyme, and it catalyzes the oxidation of this aldehyde while still in the attached state. The high-energy enzyme-substrate bond created by the oxidation is then displaced by an inorganic phosphate ion to produce a high-energy sugar-phosphate intermediate, which is thereby released from the enzyme. This intermediate then binds to the second enzyme (phosphoglycerate kinase). This enzyme catalyzes the energetically favorable transfer of the high-energy phosphate just created to ADP, forming ATP and completing the process of oxidizing an aldehyde to a carboxylic acid (see Figure 2-73).

We have shown this particular oxidation process in some detail because it provides a clear example of enzyme-mediated energy storage through coupled reactions (Figure 2-74). These reactions (steps 6 and 7) are the only ones in glycolysis that create a high-energy phosphate linkage directly from inorganic phosphate. As such, they account for the net yield of two ATP molecules and two NADH molecules per molecule of glucose (see Panel 2-8, pp. 124–125).

As we have just seen, ATP can be formed readily from ADP when reaction intermediates are formed with higher-energy phosphate bonds than those in ATP. Phosphate bonds can be ordered in energy by comparing the standard free-energy change (ΔG°) for the breakage of each bond by hydrolysis. Figure 2-75 compares the high-energy phosphoanhydride bonds in ATP with other phosphate bonds, several of which are generated during glycolysis.

Sugars and Fats Are Both Degraded to Acetyl CoA in Mitochondria

We now move on to consider stage 3 of catabolism, a process that requires abundant molecular oxygen (O₂ gas). Since the Earth is thought to have developed an atmosphere containing O₂ gas between one and two billion years ago, whereas abundant life-forms are known to have existed on the Earth for 3.5 billion years, the use of O₂ in the reactions that we discuss next is thought to be of relatively recent origin. In contrast, the mechanism used to produce ATP in Figure 2-73 does not require oxygen, and relatives of this elegant pair of coupled reactions could have arisen very early in the history of life on Earth.

In aerobic metabolism, the pyruvate produced by glycolysis is rapidly decarboxylated by a giant complex of three enzymes, called the pyruvate dehydrogenase complex. The products of pyruvate decarboxylation are a molecule of CO₂ (a waste product), a molecule of NADH, and acetyl CoA. The three-enzyme complex is located in the mitochondria of eucaryotic cells; its structure and mode of action are outlined in Figure 2-76.

The enzymes that degrade the fatty acids derived from fats likewise produce acetyl CoA in mitochondria. Each molecule of fatty acid (as the activated molecule fatty acyl CoA) is broken down completely by a cycle of reactions that trims two carbons at a time from its carboxyl end, generating one molecule of acetyl CoA for each turn of the cycle. A molecule of NADH and a molecule of FADH₂ are also produced in this process (Figure 2-77).

Sugars and fats provide the major energy sources for most non-photosynthetic organisms, including humans. However, the majority of the useful energy that can be extracted from the oxidation of both types of foodstuffs remains stored in the acetyl CoA molecules that are produced by the two types of reactions just described. The citric acid cycle of reactions, in which the acetyl group in acetyl CoA is oxidized to CO₂ and H₂O, is therefore central to the energy metabolism of aerobic organisms. In eucaryotes these reactions all take place in mitochondria, the organelle to which pyruvate and fatty acids are directed for acetyl CoA production (Figure 2-78). We should therefore not be surprised to discover that the mitochondrion is the place where most of the ATP is produced in animal cells. In contrast, aerobic bacteria carry out all of their reactions in a single compartment, the cytosol, and it is here that the citric acid cycle takes place in these cells.

The Citric Acid Cycle Generates NADH by Oxidizing Acetyl Groups to CO₂

In the nineteenth century, biologists noticed that in the absence of air (anaerobic conditions) cells produce lactic acid (for example, in muscle) or ethanol (for example, in yeast), while in its presence (aerobic conditions) they consume O₂ and produce CO₂ and H₂O. Intensive efforts to define the pathways of aerobic metabolism eventually focused on the oxidation of pyruvate and led in 1937 to the discovery of the citric acid cycle, also known as the tricarboxylic acid cycle or the Krebs cycle. The citric acid cycle accounts for about two-thirds of the total oxidation of carbon compounds in most cells, and its major end products are CO₂ and high-energy electrons in the form of NADH. The CO₂ is released as a waste product, while the high-energy electrons from NADH are passed to a membrane-bound electron-transport chain, eventually combining with O₂ to produce H₂O. Although the citric acid cycle itself does not use O₂, it requires O₂ in order to proceed because there is no other efficient way for the NADH to get rid of its electrons and thus regenerate the NAD⁺ that is needed to keep the cycle going.

The citric acid cycle, which takes place inside mitochondria in eucaryotic cells, results in the complete oxidation of the carbon atoms of the acetyl groups in acetyl CoA, converting them into CO₂. But the acetyl group is not oxidized directly. Instead, this group is transferred from acetyl CoA to a larger, four-carbon molecule, oxaloacetate, to form the six-carbon tricarboxylic acid, citric acid, for which the subsequent cycle of reactions is named. The citric acid molecule is then gradually oxidized, allowing the energy of this oxidation to be harnessed to produce energy-rich activated carrier molecules. The chain of eight reactions forms a cycle because at the end the oxaloacetate is regenerated and enters a new turn of the cycle, as shown in outline in Figure 2-79.

We have thus far discussed only one of the three types of activated carrier molecules that are produced by the citric acid cycle, the NAD⁺-NADH pair (see Figure 2-60). In addition to three molecules of NADH, each turn of the cycle also produces one molecule of FADH ₂ (reduced flavin adenine dinucleotide) from FAD and one molecule of the ribonucleotide GTP (guanosine triphosphate) from GDP. The structures of these two activated carrier molecules are illustrated in Figure 2-80. GTP is a close relative of ATP, and the transfer of its terminal phosphate group to ADP produces one ATP molecule in each cycle. Like NADH, FADH₂ is a carrier of high-energy electrons and hydrogen. As we discuss shortly, the energy that is stored in the readily transferred high-energy electrons of NADH and FADH₂ will be utilized subsequently for ATP production through the process of oxidative phosphorylation, the only step in the oxidative catabolism of foodstuffs that directly requires gaseous oxygen (O₂) from the atmosphere.

The complete citric acid cycle is presented in Panel 2-9 (pp. 126–127). The extra oxygen atoms required to make CO₂ from the acetyl groups entering the citric acid cycle are supplied not by molecular oxygen, but by water. As illustrated in the panel, three molecules of water are split in each cycle, and the oxygen atoms of some of them are ultimately used to make CO₂.

In addition to pyruvate and fatty acids, some amino acids pass from the cytosol into mitochondria, where they are also converted into acetyl CoA or one of the other intermediates of the citric acid cycle. Thus, in the eucaryotic cell, the mitochondrion is the center toward which all energy-yielding processes lead, whether they begin with sugars, fats, or proteins.

The citric acid cycle also functions as a starting point for important biosynthetic reactions by producing vital carbon-containing intermediates, such as oxaloacetate and α-ketoglutarate. Some of these substances produced by catabolism are transferred back from the mitochondrion to the cytosol, where they serve in anabolic reactions as precursors for the synthesis of many essential molecules, such as amino acids.

Electron Transport Drives the Synthesis of the Majority of the ATP in Most Cells

It is in the last step in the degradation of a food molecule that the major portion of its chemical energy is released. In this final process the electron carriers NADH and FADH₂ transfer the electrons that they have gained when oxidizing other molecules to the electron-transport chain, which is embedded in the inner membrane of the mitochondrion. As the electrons pass along this long chain of specialized electron acceptor and donor molecules, they fall to successively lower energy states. The energy that the electrons release in this process is used to pump H⁺ ions (protons) across the membrane—from the inner mitochondrial compartment to the outside (Figure 2-81). A gradient of H⁺ ions is thereby generated. This gradient serves as a source of energy, being tapped like a battery to drive a variety of energy-requiring reactions. The most prominent of these reactions is the generation of ATP by the phosphorylation of ADP.

At the end of this series of electron transfers, the electrons are passed to molecules of oxygen gas (O₂) that have diffused into the mitochondrion, which simultaneously combine with protons (H⁺) from the surrounding solution to produce molecules of water. The electrons have now reached their lowest energy level, and therefore all the available energy has been extracted from the food molecule being oxidized. This process, termed oxidative phosphorylation (Figure 2-82), also occurs in the plasma membrane of bacteria. As one of the most remarkable achievements of cellular evolution, it will be a central topic of Chapter 14.

In total, the complete oxidation of a molecule of glucose to H₂O and CO₂ is used by the cell to produce about 30 molecules of ATP. In contrast, only 2 molecules of ATP are produced per molecule of glucose by glycolysis alone.

Organisms Store Food Molecules in Special Reservoirs

All organisms need to maintain a high ATP/ADP ratio, if biological order is to be maintained in their cells. Yet animals have only periodic access to food, and plants need to survive overnight without sunlight, without the possibility of sugar production from photosynthesis. For this reason, both plants and animals convert sugars and fats to special forms for storage (Figure 2-83).

To compensate for long periods of fasting, animals store fatty acids as fat droplets composed of water-insoluble triacylglycerols, largely in specialized fat cells. And for shorter-term storage, sugar is stored as glucose subunits in the large branched polysaccharide glycogen, which is present as small granules in the cytoplasm of many cells, including liver and muscle. The synthesis and degradation of glycogen are rapidly regulated according to need. When more ATP is needed than can be generated from the food molecules taken in from the bloodstream, cells break down glycogen in a reaction that produces glucose 1-phosphate, which enters glycolysis.

Quantitatively, fat is a far more important storage form than glycogen, in part because the oxidation of a gram of fat releases about twice as much energy as the oxidation of a gram of glycogen. Moreover, glycogen differs from fat in binding a great deal of water, producing a sixfold difference in the actual mass of glycogen required to store the same amount of energy as fat. An average adult human stores enough glycogen for only about a day of normal activities but enough fat to last for nearly a month. If our main fuel reservoir had to be carried as glycogen instead of fat, body weight would need to be increased by an average of about 60 pounds.

Most of our fat is stored in adipose tissue, from which it is released into the bloodstream for other cells to utilize as needed. The need arises after a period of not eating; even a normal overnight fast results in the mobilization of fat, so that in the morning most of the acetyl CoA entering the citric acid cycle is derived from fatty acids rather than from glucose. After a meal, however, most of the acetyl CoA entering the citric acid cycle comes from glucose derived from food, and any excess glucose is used to replenish depleted glycogen stores or to synthesize fats. (While animal cells readily convert sugars to fats, they cannot convert fatty acids to sugars.)

Although plants produce NADPH and ATP by photosynthesis, this important process occurs in a specialized organelle, called a chloroplast, which is isolated from the rest of the plant cell by a membrane that is impermeable to both types of activated carrier molecules. Moreover, the plant contains many other cells—such as those in the roots—that lack chloroplasts and therefore cannot produce their own sugars or ATP. Therefore, for most of its ATP production, the plant relies on an export of sugars from its chloroplasts to the mitochondria that are located in all cells of the plant. Most of the ATP needed by the plant is synthesized in these mitochondria and exported from them to the rest of the plant cell, using exactly the same pathways for the oxidative breakdown of sugars that are utilized by nonphotosynthetic organisms (Figure 2-84).

During periods of excess photosynthetic capacity during the day, chloroplasts convert some of the sugars that they make into fats and into starch, a polymer of glucose analogous to the glycogen of animals. The fats in plants are triacylglycerols, just like the fats in animals, and differ only in the types of fatty acids that predominate. Fat and starch are both stored in the chloroplast as reservoirs to be mobilized as an energy source during periods of darkness (see Figure 2-83B).

The embryos inside plant seeds must live on stored sources of energy for a prolonged period, until they germinate to produce leaves that can harvest the energy in sunlight. For this reason plant seeds often contain especially large amounts of fats and starch—which makes them a major food source for animals, including ourselves (Figure 2-85).

Amino Acids and Nucleotides Are Part of the Nitrogen Cycle

In our discussion so far we have concentrated mainly on carbohydrate metabolism. We have not yet considered the metabolism of nitrogen or sulfur. These two elements are constituents of proteins and nucleic acids, which are the two most important classes of macromolecules in the cell and make up approximately two-thirds of its dry weight. Atoms of nitrogen and sulfur pass from compound to compound and between organisms and their environment in a series of reversible cycles.

Although molecular nitrogen is abundant in the Earth's atmosphere, nitrogen is chemically unreactive as a gas. Only a few living species are able to incorporate it into organic molecules, a process called nitrogen fixation. Nitrogen fixation occurs in certain microorganisms and by some geophysical processes, such as lightning discharge. It is essential to the biosphere as a whole, for without it life would not exist on this planet. Only a small fraction of the nitrogenous compounds in today's organisms, however, is due to fresh products of nitrogen fixation from the atmosphere. Most organic nitrogen has been in circulation for some time, passing from one living organism to another. Thus present-day nitrogen-fixing reactions can be said to perform a “topping-up” function for the total nitrogen supply.

Vertebrates receive virtually all of their nitrogen in their dietary intake of proteins and nucleic acids. In the body these macromolecules are broken down to amino acids and the components of nucleotides, and the nitrogen they contain is used to produce new proteins and nucleic acids or utilized to make other molecules. About half of the 20 amino acids found in proteins are essential amino acids for vertebrates (Figure 2-86), which means that they cannot be synthesized from other ingredients of the diet. The others can be so synthesized, using a variety of raw materials, including intermediates of the citric acid cycle as described below. The essential amino acids are made by nonvertebrate organisms, usually by long and energetically expensive pathways that have been lost in the course of vertebrate evolution.

The nucleotides needed to make RNA and DNA can be synthesized using specialized biosynthetic pathways: there are no “essential nucleotides” that must be provided in the diet. All of the nitrogens in the purine and pyrimidine bases (as well as some of the carbons) are derived from the plentiful amino acids glutamine, aspartic acid, and glycine, whereas the ribose and deoxyribose sugars are derived from glucose.

Amino acids that are not utilized in biosynthesis can be oxidized to generate metabolic energy. Most of their carbon and hydrogen atoms eventually form CO₂ or H₂O, whereas their nitrogen atoms are shuttled through various forms and eventually appear as urea, which is excreted. Each amino acid is processed differently, and a whole constellation of enzymatic reactions exists for their catabolism.

Many Biosynthetic Pathways Begin with Glycolysis or the Citric Acid Cycle

Catabolism produces both energy for the cell and the building blocks from which many other molecules of the cell are made (see Figure 2-36). Thus far, our discussions of glycolysis and the citric acid cycle have emphasized energy production, rather than the provision of the starting materials for biosynthesis. But many of the intermediates formed in these reaction pathways are also siphoned off by other enzymes that use them to produce the amino acids, nucleotides, lipids, and other small organic molecules that the cell needs. Some idea of the complexity of this process can be gathered from Figure 2-87, which illustrates some of the branches from the central catabolic reactions that lead to biosyntheses.

The existence of so many branching pathways in the cell requires that the choices at each branch be carefully regulated, as we discuss next.

Metabolism Is Organized and Regulated

One gets a sense of the intricacy of a cell as a chemical machine from the relation of glycolysis and the citric acid cycle to the other metabolic pathways sketched out in Figure 2-88. This type of chart, which was used earlier in this chapter to introduce metabolism, represents only some of the enzymatic pathways in a cell. It is obvious that our discussion of cell metabolism has dealt with only a tiny fraction of cellular chemistry.

All these reactions occur in a cell that is less than 0.1 mm in diameter, and each requires a different enzyme. As is clear from Figure 2-88, the same molecule can often be part of many different pathways. Pyruvate, for example, is a substrate for half a dozen or more different enzymes, each of which modifies it chemically in a different way. One enzyme converts pyruvate to acetyl CoA, another to oxaloacetate; a third enzyme changes pyruvate to the amino acid alanine, a fourth to lactate, and so on. All of these different pathways compete for the same pyruvate molecule, and similar competitions for thousands of other small molecules go on at the same time. A better sense of this complexity can perhaps be attained from a three-dimensional metabolic map that allows the connections between pathways to be made more directly (Figure 2-89).

The situation is further complicated in a multicellular organism. Different cell types will in general require somewhat different sets of enzymes. And different tissues make distinct contributions to the chemistry of the organism as a whole. In addition to differences in specialized products such as hormones or antibodies, there are significant differences in the “common” metabolic pathways among various types of cells in the same organism.

Although virtually all cells contain the enzymes of glycolysis, the citric acid cycle, lipid synthesis and breakdown, and amino acid metabolism, the levels of these processes required in different tissues are not the same. For example, nerve cells, which are probably the most fastidious cells in the body, maintain almost no reserves of glycogen or fatty acids and rely almost entirely on a constant supply of glucose from the bloodstream. In contrast, liver cells supply glucose to actively contracting muscle cells and recycle the lactic acid produced by muscle cells back into glucose (Figure 2-90). All types of cells have their distinctive metabolic traits, and they cooperate extensively in the normal state, as well as in response to stress and starvation. One might think that the whole system would need to be so finely balanced that any minor upset, such as a temporary change in dietary intake, would be disastrous.

In fact, the metabolic balance of a cell is amazingly stable. Whenever the balance is perturbed, the cell reacts so as to restore the initial state. The cell can adapt and continue to function during starvation or disease. Mutations of many kinds can damage or even eliminate particular reaction pathways, and yet—provided that certain minimum requirements are met—the cell survives. It does so because an elaborate network of control mechanisms regulates and coordinates the rates of all of its reactions. These controls rest, ultimately, on the remarkable abilities of proteins to change their shape and their chemistry in response to changes in their immediate environment. The principles that underlie how large molecules such as proteins are built and the chemistry behind their regulation will be our next concern.

Summary

Glucose and other food molecules are broken down by controlled stepwise oxidation to provide chemical energy in the form of ATP and NADH. These are three main sets of reactions that act in series—the products of each being the starting material for the next: glycolysis (which occurs in the cytosol), the citric acid cycle (in the mitochondrial matrix), and oxidative phosphorylation (on the inner mitochondrial membrane). The intermediate products of glycolysis and the citric acid cycle are used both as sources of metabolic energy and to produce many of the small molecules used as the raw materials for biosynthesis. Cells store sugar molecules as glycogen in animals and starch in plants; both plants and animals also use fats extensively as a food store. These storage materials in turn serve as a major source of food for humans, along with the proteins that comprise the majority of the dry mass of the cells we eat.