Life is possible because molecules in biological systems frequently undergo chemical reactions, enabling organisms to replicate DNA, synthesize RNA and protein molecules, pump small molecules into and out of cells, and use energy in the form of food or light. In this section we discuss the physical principles governing chemical reactions. We review the fundamental laws of thermodynamics and the roles of catalysts in accelerating the rates of reactions between biomolecules. Finally, we describe how energy, in the form of high-energy phosphate compounds, is harnessed to drive certain chemical reactions that would otherwise occur too rarely or too slowly to be useful to living systems.

Chemical reactions involve the breakage of covalent bonds and the formation of new bonds. Typically, chemical reactions are written with the reactants, or starting molecules, on the left and the products on the right, connected by an arrow indicating the direction of the reaction. For example, the expression

describes a very common reaction in biology: the breakage of a phosphorus–oxygen bond in the nucleotide adenosine triphosphate (ATP) to produce adenosine diphosphate (ADP) and inorganic phosphate (P_i). A more detailed representation of this reaction can be drawn to indicate, with curved arrows, the direction in which electrons are moving (Figure 3-26).

Reactions of this type involve the attack of a nucleophile, a strongly electronegative atom, such as oxygen or nitrogen, on a less electronegative atom, such as phosphorus or carbon. When the nucleophile initiating the reaction is part of a water molecule, as in Figure 3-26, the reaction is known as hydrolysis.

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Most chemical reactions that take place in biological systems are not spontaneous; if they were, they would be impossible to control, and life as we know it could not exist. Instead, the starting and ending points of chemical reactions are bridged by an energy barrier, called the activation energy, that separates the reactants from the products (Figure 3-27). As reactants come together and bonds are breaking and forming, the reacting species progress through a high-energy state called a transition state. Once past this state, the reaction proceeds spontaneously because it is energetically favorable. Any chemical reaction can, in principle, proceed in the forward or reverse direction—toward products or reactants. In practice, however, most reactions characteristically tend to proceed more rapidly in one direction than the other, due to differences in the relative energetic stability of products versus reactants. In Figure 3-27, the difference in energy between the products and the transition state is greater than the difference in energy between the reactants and the transition state. Hence, at equilibrium, there will be more products than reactants, because the energetic barrier to the reverse reaction is higher. (This is further discussed later in the chapter.)

A reaction mechanism is the sequence of individual steps that take place during the conversion of reactants to products. It shows which bonds are breaking and forming and which species are reaction intermediates—substances that form and exist for an extremely short time before being converted to other intermediates or to reaction products. Understanding the mechanisms of chemical reactions that occur in living systems is important in molecular biology because it helps us understand how these reactions are used and controlled during such processes as cell growth and responses to chemical stimuli.

Although it is often difficult to determine reaction mechanisms, understanding the reaction kinetics, the rates at which the reaction proceeds in the forward and reverse directions, can provide important clues. Reaction rates are a function of the concentration of reactants (when the reaction is catalyzed by an enzyme, reactants are called substrates) and their tendency to react. Chemical reactions require collisions between molecules, and collisions occur more frequently when there are more molecules per unit volume. It is important to note that reaction rates also depend on the individual steps that make up the reaction. Usually, one step is slower than the others, and this slowest step, the rate-limiting step, governs the overall reaction rate. For example, suppose the reaction A + 2B → 2C is found, by experiment, to occur with a reaction rate proportional to the concentration of A multiplied by the concentration of B. We can write this as:

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where k is a value called the rate constant, a property of the overall reaction that quantitatively describes the tendency to react. A possible mechanism for this reaction might be:

Step 1, the rate-limiting step, produces an intermediate, D, that rapidly reacts with B in step 2 to yield product C. This reaction mechanism is not the only one consistent with the observed reaction kinetics, but it provides a hypothesis that could be tested experimentally.

Many biologically important reactions involve not just two but many individual steps. As we shall see, the identification of so many individual steps can be extremely challenging.

Living systems demand an almost constant input of energy, and as a result, organisms devote considerable molecular machinery to obtaining and using it. Biology obeys the physical laws of thermodynamics, which are the foundation for understanding energy and its effects on matter. In thermodynamics, a system is defined as a container or organism or other portion of the universe that is under study, and the rest of the universe, outside the system of interest, is called the surroundings.

The first law of thermodynamics states that energy can never be created or destroyed; in other words, the energy of a system is conserved. In a thermodynamic system, all readily occurring (spontaneous) processes take place without the input of additional energy from outside. The first law of thermodynamics cannot predict whether a process is spontaneous, however. For example, heat spontaneously transfers from a warmer object to a cooler one, never the reverse. Yet transfer in either direction is consistent with the first law of thermodynamics, because the total energy of the system remains unchanged in either case. To determine which direction of a process or reaction is spontaneous, we need additional criteria.

According to the second law of thermodynamics, all spontaneous processes take place with an increase in disorder, or entropy, of the system. Consider, for instance, two vessels of equal volume, one filled with plain water and the other with water containing a dye. When a connecting valve is opened, the dye molecules become randomly but equally distributed between the two vessels (Figure 3-28). The likelihood of all the dye molecules spontaneously remaining in the first vessel, or all the dye molecules moving into the other vessel, is essentially zero, even though the energy of either of these arrangements is no different from that of the evenly distributed molecules. The spontaneity of a process, such as a biochemical reaction, cannot be predicted from knowledge of the system’s entropy change alone.

Every closed system, where no reactants or products can escape, tends toward equilibrium, a state in which the forward and reverse reaction rates are exactly balanced. The approach to equilibrium is accompanied by a transfer of energy from one form to another, such as from potential (stored) energy to kinetic energy. The concept of free energy (G) provides a useful way to express this energy change. Free energy, denoted by the symbol G (for nineteenth-century physicist Josiah Willard Gibbs), is energy that is available to do work. The second law of thermodynamics states that free energy always decreases—that is, the free-energy change (ΔG) is negative—in spontaneous reactions that occur without a change in temperature or pressure; at equilibrium, ΔG is zero. Free energy lost during the approach to equilibrium is either converted into heat or used to increase the amount of entropy.

The tendency of a chemical reaction to proceed to completion can be expressed as an equilibrium constant, which is related to the standard Gibbs free-energy change (ΔG^o) of the reaction by the expression:

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where R is the universal gas constant (1.987 cal/mol • K),T is the absolute temperature (298 at 25°C), and ln K_eq is the natural logarithm of the equilibrium constant K_eq. The standard state (under which K_eq is measured) requires that all reactants be present at a concentration of 1 mol/L (1 m). For any compound, the ΔG^o of formation ( ) is the change in free energy that accompanies the formation of 1 mol of that substance from its component elements in their standard states (the most stable form of each element at 25°C and 100 kPa (kilopascals) of atmospheric pressure). Because K_eq can be measured experimentally, this relationship gives us a way of calculating ΔG^o—the thermodynamic constant characteristic of each reaction. When K_eq is much greater than 1, ΔG^o is large and negative; reactions of this nature tend to go to completion. In contrast, when K_eq is much smaller than 1, ΔG^o is large and positive; reactions with this property are not spontaneous and require energy to drive them to completion.

The ΔG of a chemical reaction reflects its equilibrium. However, ΔG, whether positive or negative, bears no relationship to the rate of the reaction. For example, the series of reactions that convert glucose to H₂O and CO₂ in most cells is thermodynamically favorable, with a net ΔG that is highly negative. However, a large, negative ΔG does not correlate with a fast reaction. A sterile solution of glucose will sit on a laboratory shelf for months with no detectable change. Glucose remains stable because of the substantial activation energy barrier that exists for its reaction. Virtually every chemical reaction in a cell occurs at a significant rate only because of the action of enzymes, which, like all catalysts, are molecules that dramatically increase the rate of specific chemical reactions without being consumed in the process.

Catalysts function by lowering the activation energy for a particular reaction without affecting the reaction equilibrium. Given that a catalyst changes the reaction’s rate but not its equilibrium, it must change the rate of the reverse reaction to the same extent as the rate of the forward reaction. Catalysts can do this because they enable the reaction to proceed by a different mechanism than that of the uncatalyzed reaction. For example, enzymes bind to the transition state of the reactants by providing a molecular surface complementary to its shape and charge. Because of this favorable interaction, binding of an enzyme stabilizes the transition state, reducing the activation energy for the reaction and thus greatly enhancing the reaction rate. Additional contributions to catalysis occur when reacting molecules—substrates—bind to an enzyme in an orientation that favors the reaction and when chemical groups in the enzyme bind metal ions or protons that participate in the reaction. As a consequence of these effects, enzymes often increase reaction rates 10¹²-fold or more above the rate of the uncatalyzed reaction.

Most cellular enzymes are proteins, though some RNA molecules also have catalytic activity. In general, each enzyme catalyzes a specific reaction, and each reaction in a cell is catalyzed by a different enzyme. Thus, many thousands of enzymes are required in each cell. Because enzymes are exquisitely capable of discriminating between substrates, and because they are subject to various regulatory mechanisms, cells can selectively adjust reaction rates. Such selectivity is critical for the effective control of cellular processes. By enabling specific reactions to occur at particular times and locations within a cell or organism, enzymes determine how chemicals and energy are channeled into biological activities. Enzyme function is described in detail in Chapter 5.

The formation and breakdown of adenosine triphosphate (ATP) (and in some cases guanosine triphosphate, GTP) links the molecule-making and molecule-degrading pathways of cellular metabolism. The formation of this critical energy-storing molecule from inorganic phosphate and adenosine diphosphate (ADP), by the creation of a phosphodiester linkage, is coupled to some of the steps of degradative metabolism and, in plants, to photosynthesis.

Because the hydrolysis of a phosphodiester bond is exothermic under physiological conditions, energy is released when ATP is hydrolyzed to ADP. In turn, the free energy stored in the phosphodiester bonds of ATP is used to drive biosynthetic reactions of metabolism. Although energy is required for bond breaking in ATP, the products of the reaction (ADP and phosphate) form highly favorable interactions with water. Thus, hydration of the breakdown products of ATP more than makes up for the input energy necessary to break the bond in the first place, resulting in an overall energetically favorable process. Almost as soon as it is formed in concert with a coupled degradative reaction (or by photosynthesis), ATP is consumed by enzymes to provide the energy necessary to propel another reaction to completion. In this way, ATP functions as a transient vehicle of intracellular energy transfer (Highlight 3-2).

ATP consists of an adenosine nucleoside (base + ribose) and three phosphate groups (see Figure 3-26). The phosphate groups, starting with the one directly bonded to ribose, are referred to as the alpha (α), beta (β), and gamma (γ) phosphates, respectively. Like other such high-energy molecules, ATP contains bonds—in this case, the phosphodiester bonds between the phosphate groups—that undergo breakdown by water (hydrolysis) to release significant free energy. The second and third (β and γ) phosphate groups of ATP are unusually rich in chemical energy: the net change in energy (ΔG) on hydrolysis of ATP to produce ADP and inorganic phosphate (P_i) inside a living cell is −12 kcal/mol. This large negative change in free energy makes the breakdown of ATP thermodynamically favorable and hence ATP is valuable for chemically storing energy that can be used to do work. The stored energy is captured when the cleaved phosphate group is transferred to another small molecule or protein as part of a metabolic pathway. Many enzymes harness ATP hydrolysis in this way to perform the work of the cell.

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Note that a single hydrolytic reaction produces one P_i and ADP, and the ADP can be broken down further to yield a second P_i and adenosine monophosphate (AMP). But ATP can also be hydrolyzed to AMP directly, with the release of inorganic diphosphate, or pyrophosphate (PP_i). This latter reaction is effectively irreversible in the cellular environment, because enzymes called pyrophosphatases rapidly hydrolyze pyrophosphate into two phosphates. As a result, it would be very difficult to accumulate sufficient concentrations of pyrophosphate in the cell to drive the reaction in the reverse direction. Thus, the hydrolysis of ATP to AMP can be used to render a coupled process effectively irreversible.

For example, DNA and RNA are synthesized from nucleoside triphosphate precursors through a reaction that forms a phosphodiester bond and releases PP_i. The required free energy of bond formation comes in part from the concomitant splitting of the high-energy pyrophosphate group (by pyrophosphatase) as it is released. Interestingly, when no pyrophosphatase is present, such as in a laboratory test tube, robust DNA and RNA synthesis can still occur in this reaction. This is because the thermodynamics of phosphate bond formation and cleavage is only part of the story; base stacking and base-pair formation in polynucleotides are energetically favorable and therefore contribute some of the push toward polynucleotide synthesis. During protein synthesis in cells, amino acids are activated for peptide bond formation by linkage to AMP with the release of PP_i. Again, the splitting of the pyrophosphate group helps drive the reaction irreversibly toward formation of the activated amino acid–AMP (aminoacyl adenylate), the substrate for protein synthesis.

In addition to phosphorus–oxygen bonds, phosphorus–nitrogen and sulfur–carbon bonds also release significant free energy on hydrolysis, and these bonds are found in other important classes of energy-storing compounds that are used to drive reactions in biology. For example, the sulfur–carbon (thioester) bond in acetyl-coenzyme A is the primary source of chemical energy for fatty acid biosynthesis. The free energy released by hydrolysis of high-energy bonds ranges in value, and its utility comes from the coupling of the released energy to another reaction to drive it forward. Thus, the coupling of biosynthetic reactions having a positive (unfavorable) ΔG value with the breakage of high-energy bonds that has a negative ΔG of greater absolute value ensures that the equilibrium favors synthesis of the biomolecule over its breakdown (Figure 3-29).

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The overall free-energy change associated with a series of linked reactions determines whether a particular reaction in the series will occur. Reactions with small positive ΔG values, which in isolation would be unfavorable and would not take place, are often embedded in metabolic pathways in which they precede reactions having large negative ΔG values. It is critical to bear in mind that no single biochemical reaction, and no single pathway, takes place in isolation. Instead, the overall equilibria of the huge network of pathways within a cell are constantly adjusting to changing concentrations of substrates as the cell grows and responds to its environment.

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