CHAPTER 13 INTRODUCTION TO METABOLISM sea fish, into light. Photosynthetic organisms transduce light energy into all these other forms of energy. The chemical mechanisms that underlie biological energy transductions have fascinated and challenged biologists for centuries. The French chemist Antoine Lavoisier recognized that animals somehow transform chemical fuels (foods) into heat and that this process of respiration is essential to life. He observed that in general, respiration is nothing but a slow combustion of carbon and hydrogen, which is entirely similar to that which occurs in a lighted lamp or candle, and that, from this point of view, animals that respire are true combustible bodies that burn and consume themselves…. One may say that this analogy between combustion and respiration has not escaped the notice of the poets, or rather the philosophers of antiquity, and which they had expounded and interpreted. This fire stolen from heaven, this torch of Prometheus, does not only represent an ingenious and poetic idea, it is a faithful picture of the operations of nature, at least for animals that breathe; one may therefore say, with the ancients, that the torch of life lights itself at the moment the infant breathes for the first time, and it does not extinguish itself except at death.i A portrait by Jacques Louis David of Antoine Lavoisier (1743–1794) in the laboratory with chemist Marie Anne Pierrette Paulze (1758–1836), his wife. We now understand much of the chemistry underlying that “torch of life.” Biological energy transductions obey the same chemical and physical laws that govern all other natural processes, and many of the types of chemical reactions that occur in living organisms have been long known to organic chemists. One unique feature of cellular chemistry is its exquisitely sensitive regulation by a variety of mechanisms that respond to changes in the external and internal circumstances of the cell and organism. In this chapter we lay out the foundational principles for understanding the reactions of metabolism that follow in Part II. We first review the laws of thermodynamics and the quantitative relationships among free energy, enthalpy, and entropy. We then review the common types of biochemical reactions that occur in living cells, reactions that harness, store, transfer, and release the energy taken up by organisms from their surroundings. Our focus then shi s to reactions that have special roles in biological energy exchanges, particularly those involving the cofactors ATP (for phosphoryl transfers) and NADH (for electron transfers). Finally, we look at the most common of the strategies for regulating biochemical reactions. Watch for examples of these principles as you read this chapter: The chemical changes and energy transductions in living organisms follow the laws of thermodynamics. The free-energy change is the maximum energy made available to do work when a chemical reaction occurs. If two reactions can be combined to yield a third reaction, the overall free energy change is the sum of the two. Cells accomplish energy-requiring chemical work by coupling an energy- releasing (exergonic) reaction such as the cleavage of ATP to an endergonic reaction (which requires energy input). Although thousands of different chemical reactions occur in the biosphere, most of them fall within a small set of reaction types. ATP is the universal energy currency in living organisms. Transfer of its phosphoryl group to a water molecule or metabolic intermediates provides the energetic push for muscle contraction, the pumping of solutes against concentration gradients, and the synthesis of complex molecules. Oxidation-reduction reactions indirectly provide much of the energy needed to make ATP. Reduced substrates such as glucose are oxidized in several steps, with the energy of oxidation steps conserved in the form of a reduced cofactor, NADH. Energy stored in NADH is used to drive the synthesis of ATP. To respond to changes in external circumstances, cells must regulate enzyme activities, by changing either the number of enzyme molecules or the catalytic activity of preexisting enzyme molecules. From a memoir by Armand Seguin and Antoine Lavoisier, dated 1789, quoted in A. Lavoisier, Oeuvres de Lavoisier, Imprimerie Impériale, Paris, 1862. i 13.1 Bioenergetics and Thermodynamics Bioenergetics is the quantitative study of energy transductions — changes of one form of energy into another — that occur in living cells, and of the nature and function of the chemical processes underlying these transductions. Although many of the principles of thermodynamics have been introduced in earlier chapters and may be familiar to you, a review of the quantitative aspects of these principles is useful here. Biological Energy Transformations Obey the Laws of Thermodynamics Many quantitative observations made by physicists and chemists on the interconversion of different forms of energy led, in the nineteenth century, to the formulation of two fundamental laws of thermodynamics. The first law is the principle of the conservation of energy: for any physical or chemical change, the total amount of energy in the universe remains constant; energy may change form or it may be transported from one region to another, but it cannot be created or destroyed. The second law of thermodynamics, which can be stated in several forms, says that the universe always tends toward increasing disorder: in all natural processes, the entropy of the universe increases. Living organisms consist of collections of molecules much more highly organized than the surrounding materials from which they are constructed, and organisms maintain and produce order, seemingly immune to the second law of thermodynamics. But living organisms do not violate the second law; they operate strictly within it. To discuss the application of the second law to biological systems, we must first define those systems and their surroundings. The reacting system is the collection of matter that is undergoing a particular chemical or physical process; it may be an organism, a cell, or two reacting compounds. The reacting system and its surroundings together constitute the universe. In the laboratory, some chemical or physical processes can be carried out in isolated or closed systems, in which no material or energy is exchanged with the surroundings. Living cells and organisms, however, are open systems, exchanging both material and energy with their surroundings; living systems are never at equilibrium with their surroundings, and the constant transactions between system and surroundings explain how organisms can create order within themselves while operating within the second law of thermodynamics. In Chapter 1 (p. 21) we defined three thermodynamic quantities that describe the energy changes occurring in a chemical reaction: Free energy, G (for J. Willard Gibbs), expresses the amount of energy capable of doing work during a reaction at constant temperature and pressure. When a reaction proceeds with the release of free energy (that is, when the system changes so as to possess less free energy), the free- energy change, ∆G, has a negative value and the reaction is said to be exergonic. In endergonic reactions, the system gains free energy and ∆G is positive. Enthalpy, H, is the heat content of the reacting system. It reflects the number and kinds of chemical bonds (covalent and noncovalent) in the reactants and products. When a chemical reaction releases heat, it is said to be exothermic; the heat content of the products is less than that of the reactants, and the change in enthalpy, ∆H, has, by convention, a negative value. Reacting systems that take up heat from their surroundings are endothermic and have positive values of ∆H. Entropy, S, is a quantitative expression for the randomness or disorder in a system (see Box 1-3). When the products of a reacting system are less complex and more disordered than the reactants, the reaction is said to proceed with a gain in entropy. The units of ∆G and ∆H are joules/mole or calories/mole (recall that 1 cal =4.184 J); units of entropy are joules/mole • Kelvin (J /mol∙K) (Table 13-1). TABLE 13-1 Some Physical Constants and Units Used in Thermodynamics Boltzmann constant, k= 1.381× 10−23 J /K Avogadro’s number, N= 6.022× 1023 mol−1 Faraday constant, F= 96,480 J /V ∙mol Gas constant, R= 8.315 J /mol∙K (= 1.987 cal/mol∙K) Units of ∆G and ∆H are J/mol (or cal/mol) Units of ∆S are J/mol • K (or cal/mol • K) 1 cal = 4.184 J Units of absolute temperature, T, are Kelvin, K 25 °C = 298 K At 25 °C, RT= 2.478 kJ/mol (= 0.592 kcal/mol) Under the conditions existing in biological systems (including constant temperature and pressure), changes in free energy, enthalpy, and entropy are related to each other quantitatively by the equation ΔG= ΔH−TΔS (13-1) in which ∆G is the change in Gibbs free energy of the reacting system, ∆H is the change in enthalpy of the system, T is the absolute temperature, and ∆S is the change in entropy of the system. By convention, ∆S has a positive sign when entropy increases and ∆H, as noted above, has a negative sign when heat is released by the system to its surroundings. Either of these conditions, both of which are typical of energetically favorable processes, tends to make ∆G negative. In fact, ∆G of a spontaneously reacting system is always negative. The second law of thermodynamics states that the entropy of the universe increases during all chemical and physical processes, but it does not require that the entropy increase take place in the reacting system itself. The order produced within cells as they grow and divide is more than compensated for by the disorder they create in their surroundings in the course of growth and division (see Box 1-3, case 2). In short, living organisms preserve their internal order by taking from their surroundings free energy in the form of nutrients or sunlight, and returning to their surroundings an equal amount of energy as heat and entropy. Cells are isothermal systems — they function at essentially constant temperature (and also function at constant pressure). Heat flow is not a source of energy for cells, because heat can do work only as it passes to a zone or an object at a lower temperature. The energy that cells can and must use is free energy, described by the Gibbs free-energy function G, which allows prediction of the direction of chemical reactions, their exact equilibrium position, and the amount of work they can (in theory) perform at constant temperature and pressure. Heterotrophic cells acquire free energy from nutrient molecules, and photosynthetic cells acquire it from absorbed solar radiation. Both kinds of cells transform this free energy into ATP and other energy-rich compounds capable of providing energy for biological work at constant temperature. Standard Free-Energy Change Is Directly Related to the Equilibrium Constant The composition of a reacting system (a mixture of chemical reactants and products) tends to continue changing until equilibrium is reached. (In the case of an organism, equilibrium is reached only a er death and complete decay.) At the equilibrium concentration of reactants and products, the rates of the forward and reverse reactions are exactly equal and no further net change occurs in the system. The concentrations of reactants and products at equilibrium define the equilibrium constant, Keq (p. 23). In the general reaction aA +bB ⇌ cC +dD where a, b, c, and d are the number of molecules of A, B, C, and D participating, the equilibrium constant is given by Keq = (13-2) where [A]eq, [B]eq, [C]eq, and [D ]eq are the molar concentrations of the reaction components at the point of equilibrium. When a reacting system is not at equilibrium, the tendency to move toward equilibrium represents a driving force, the magnitude of which can be expressed as the free-energy change for the reaction, ∆G. Under standard conditions of temperature and pressure and when reactants and products are initially present at 1 M concentrations or, for gases, at partial pressures of 101.3 kilopascals (kPa), or 1 atm, the force driving the system toward equilibrium is defined as the standard free-energy change, ΔG°. By this definition, the standard state for reactions that involve hydrogen ions is [H+]= 1 M , or pH 0. Most biochemical reactions, however, occur in well-buffered aqueous solutions near pH 7; both the pH and the concentration of water (55.5 M) are essentially constant. KEY CONVENTION [C]ceq[D ]deq [A]aeq[B]beq For convenience of calculations, biochemists define a different standard state from that used in chemistry and physics: in the biochemical standard state, [H+] is 10−7M (pH 7) and [H2O] is 55.5 M. For reactions that involve M g2+ (which include most of those with ATP as a reactant), [M g2+] in solution is commonly taken to be constant at 1 mM. Physical constants based on this biochemical standard state are called standard transformed constants and are written with a prime (such as ΔG′° and K′eq) to distinguish them from the untransformed constants used by chemists and physicists. (Note that most other textbooks use the symbol ΔG°′ rather than ΔG′°. Our use of ΔG′°, recommended by an international committee of chemists and biochemists, is intended to emphasize that the transformed free-energy change, ΔG′°, is the criterion for equilibrium.) For simplicity, we will herea er refer to these transformed constants as standard free-energy changes and standard equilibrium constants. KEY CONVENTION In another simplifying convention used by biochemists, when H2O, H+, and/or M g2+ are reactants or products, their concentrations are not included in equations such as Equation 13- 2 but are instead incorporated into the constants K′eq and ΔG′°.
Just as K′eq is a physical constant characteristic for each reaction, so too is ΔG′° a constant. As we noted in Chapter 6, there is a simple relationship between K′eq and ΔG′°: ΔG′°= −RTlnK′eq (13-3) The standard free-energy change of a chemical reaction is simply an alternative mathematical way of expressing its equilibrium constant. Table 13-2 shows the relationship between ΔG′° and K′eq. If the equilibrium constant for a given chemical reaction is 1.0, the standard free-energy change of that reaction is 0.0 (the natural logarithm of 1.0 is zero). If K′eq of a reaction is greater than 1.0, its ΔG′° is negative. If K′eq is less than 1.0, ΔG′° is positive. Because the relationship between ΔG′° and K′eq is exponential, relatively small changes in ΔG′° correspond to large changes in K′eq. TABLE 13-2 Relationship between Equilibrium Constants and Standard Free-Energy Changes of Chemical Reactions ΔG′° K′eq (kJ/mol) (kcal/mol) 103 −17.1 −4.1 102 −11.4 −2.7 101 −5.7 −1.4 a 1 0.0 0.0 10−1 5.7 1.4 10−2 11.4 2.7 10−3 17.1 4.1 10−4 22.8 5.5 10−5 28.5 6.8 10−6 34.2 8.2 Although joules and kilojoules are the standard units of energy and are used throughout this text, biochemists and nutritionists sometimes express ΔG′° values in kilocalories per mole. We have therefore included values in both kilojoules and kilocalories in this table and in Tables 13-4 and 13-6. To convert kilojoules to kilocalories, divide the number of kilojoules by 4.184. It may be helpful to think of the standard free-energy change in another way. ΔG′° is the difference between the free-energy content of the products and the free-energy content of the reactants, under standard conditions. When ΔG′° is negative, the products contain less free energy than the reactants and the reaction will proceed spontaneously under standard conditions; all chemical reactions tend to go in the direction that results in a decrease in the free energy of the system. A positive value of ΔG′° means that the products of the reaction contain more free energy than the reactants, and this reaction will tend to go in the reverse direction if we start with 1.0 M concentrations of all components (standard conditions). Table 13-3 summarizes these points. a TABLE 13-3 Relationships among K′eq, ΔG′°, and the Direction of Chemical Reactions When K′eq is … ΔG′° is … Starting with all components at 1 , the reaction … >1.0 negative proceeds forward 1.0 zero is at equilibrium <1.0 positive proceeds in reverse WORKED EXAMPLE 13-1 Calculation of ΔG′° Calculate the standard free-energy change of the reaction catalyzed by the enzyme phosphoglucomutase, G lucose 1-phosphate ⇌ glucose 6-phosphate given that, starting with 20 mM glucose 1-phosphate and no glucose 6-phosphate, the final equilibrium mixture at 25 °C and pH 7.0 contains 1.0 mM glucose 1-phosphate and 19 mM glucose 6- phosphate. Does the reaction in the direction of glucose 6- phosphate formation proceed with a loss or a gain of free energy? SOLUTION: First we calculate the equilibrium constant: K′eq = = = 19 We can now calculate the standard free-energy change: ΔG′° = −RT ln K′eq = −(8.315 J /mol∙K)(298 K)(ln 19) = −7.3 kJ /mol Because the standard free-energy change is negative, the conversion of glucose 1-phosphate to glucose 6-phosphate proceeds with a loss (release) of free energy. (For the reverse reaction, ΔG′° has the same magnitude but the opposite sign.) Table 13-4 gives the standard free-energy changes for some representative chemical reactions. Note that hydrolysis of simple esters, amides, peptides, and glycosides, as well as rearrangements and eliminations, proceed with relatively small standard free-energy changes, whereas hydrolysis of acid anhydrides is accompanied by relatively large decreases in standard free energy. The complete oxidation of organic compounds such as glucose or palmitate to CO2 and H2O, which in cells requires many steps, results in very large decreases in standard free energy. However, standard free-energy changes such as those in Table 13-4 indicate how much free energy is [glucose 6-phosphate]eq [glucose 1-phosphate]eq 19 mM 1.0 mM available from a reaction under standard conditions. To describe the energy released under the conditions existing in cells, an expression for the actual free-energy change is essential. TABLE 13-4 Standard Free-Energy Changes of Some Chemical Reactions ΔG′° Reaction type (kJ/mol) (kcal/mol) Hydrolysis reactions Acid anhydrides Acetic anhydride+ H2O → 2 acetate −91.1 −21.8 ATP + H2O → AD P + Pi −30.5 −7.3 ATP + H2O → AM P + PPi −45.6 −10.9 PPi+ H2O → 2Pi −19.2 −4.6 U D P-glucose+ H2O → U M P + glucose 1-phosphate −43.0 −10.3 Esters Ethyl acetate+ H2O → ethanol+ acetate −19.6 −4.7 G lucose 6-phosphate+ H2O → glucose+ Pi −13.8 −3.3 Amides and peptides G lutamine+ H2O → glutamate+ NH+4 −14.2 −3.4 G lycylglycine+ H2O → 2 glycine −9.2 −2.2 Glycosides M altose+ H2O → 2 glucose −15.5 −3.7 Lactose+ H2O → glucose+ galactose −15.9 −3.8 Rearrangements G lucose 1-phosphate→ glucose 6-phosphate −7.3 −1.7 Fructose 6-phosphate→ glucose 6-phosphate −1.7 −0.4 Elimination of water M alate→ fumarate+ H2O 3.1 0.8 Oxidations with molecular oxygen G lucose+ 6O2 → 6CO2+ 6H2O −2,840 −686 Palmitate+ 23O2 → 16CO2+ 16H2O −9,770 −2,338 Actual Free-Energy Changes Depend on Reactant and Product Concentrations We must be careful to distinguish between two different quantities: the actual free-energy change, ∆G, and the standard free-energy change, ΔG′°. Each chemical reaction has a characteristic standard free-energy change, which may be positive, negative, or zero, depending on the equilibrium constant of the reaction. The standard free-energy change tells us in which direction and how far a given reaction must go to reach equilibrium when the initial concentration of each component is 1.0 M, the pH is 7.0, the temperature is 25 °C, and the pressure is 101.3 kPa (1 atm). Thus ΔG′° is a constant: it has a characteristic, unchanging value for a given reaction. But the actual free-energy change, ΔG, is a function of reactant and product concentrations and of the temperature prevailing during the reaction, none of which will necessarily match the standard conditions as defined above. Moreover, the ΔG of any reaction proceeding spontaneously toward its equilibrium is always negative, becomes less negative as the reaction proceeds, and is zero at the point of equilibrium, indicating that no more work can be done by the reaction. ΔG and ΔG′° for any reaction aA +bB ⇌ cC +dD are related by the equation (13-4) in which the terms in red are those actually prevailing in the system under observation. The concentration terms in this equation express the effects commonly called mass action, and the term [C]c[D ]d/[A]a[B]b is called the mass-action ratio, Q. Thus Equation 13-4 can be expressed as ΔG= ΔG′°+RT ln Q. As an example, let us suppose that the reaction A + B ⇌ C + D is taking place under the standard conditions of temperature (25 °C) and pressure (101.3 kPa) but that the concentrations of A, B, C, and D are not equal and none of the components is present at the standard concentration of 1.0 M. To determine the actual free- energy change, ΔG, under these nonstandard conditions of concentration as the reaction proceeds from le to right, we simply enter the actual concentrations of A, B, C, and D in Equation 13-4; the values of R, T, and ΔG′° are the standard values. ΔG is negative and approaches zero as the reaction proceeds, because the actual concentrations of A and B decrease and the concentrations of C and D increase. Notice that when a reaction is at equilibrium — when there is no force driving the reaction in either direction and ΔG is zero — Equation 13-4 reduces to 0= ΔG= ΔG′°+RT ln or ΔG′°= −RT ln K′eq which is the equation relating the standard free-energy change and equilibrium constant (Eqn 13-3). The criterion for spontaneity of a reaction is the value of ΔG, not ΔG′°. A reaction with a positive ΔG′° can go in the forward [C]eq[D ]eq [A]eq[B]eq direction if ΔG is negative. This is possible if the term RT ln ([products]/[reactants]) in Equation 13-4 is negative and has a larger absolute value than ΔG′°. For example, the immediate removal of the products of a reaction by an enzyme that degrades the product can keep the ratio [products]/[reactants] well below 1, such that the term RT ln ([products]/[reactants]) has a large, negative value. This is a quantitative expression of Le Chatelier’s principle. ΔG′° and ΔG are expressions of the maximum amount of free energy that a given reaction can theoretically deliver — an amount of energy that could be realized only if a perfectly efficient device were available to trap or harness it. Given that no such device is possible (some energy is always lost to entropy during any process), the amount of work done by the reaction at constant temperature and pressure is always less than the theoretical amount. Another important point is that some thermodynamically favorable reactions (that is, reactions for which ΔG′° is large and negative) do not occur at measurable rates. For example, combustion of firewood to CO2 and H2O is very favorable thermodynamically, but firewood remains stable for years because the activation energy (see Figs. 6-2, 6-3) for the combustion reaction is higher than the energy available at room temperature. If the necessary activation energy is provided (with a lighted match, for example), combustion will begin, converting the wood to the more stable products CO2 and H2O and releasing energy as heat and light. The heat released by this exothermic reaction provides the activation energy for combustion of neighboring regions of the firewood; the process is self- perpetuating. Thermodynamics allows us to predict which direction a process will tend to go; how fast it will go is the subject of kinetics. In living cells, reactions that would be extremely slow if uncatalyzed are caused to proceed not by supplying additional heat but by lowering the activation energy through use of an enzyme catalyst. An enzyme provides an alternative reaction pathway with a lower activation energy than the uncatalyzed reaction, so that at body temperature a large fraction of the substrate molecules have enough thermal energy to overcome the activation barrier, and the reaction rate increases dramatically. The free-energy change for a reaction is independent of the pathway by which the reaction occurs; it depends only on the nature and concentration of the initial reactants and the final products. Enzymes cannot, therefore, change equilibrium constants; but they can and do increase the rate at which a reaction proceeds in the direction dictated by thermodynamics (see Section 6.2). Standard Free-Energy Changes Are Additive In the case of two sequential chemical reactions, A ⇌ B and B ⇌ C, each reaction has its own equilibrium constant and each has its characteristic standard free-energy change, ΔG′1° and ΔG′2°. As the two reactions are sequential, B cancels out to give the overall reaction A ⇌ C, which has its own equilibrium constant and thus its own standard free-energy change, ΔG′°Sum. The ΔG′°values of sequential chemical reactions are additive. For the overall reaction A ⇌ C, ΔG′°Sum is the sum of the individual standard free-energy changes, ΔG′1° and ΔG′2°, of the two reactions: ΔG′°Sum = ΔG′1°+ΔG′2°. This principle of bioenergetics explains how a thermodynamically unfavorable (endergonic) reaction can be driven in the forward direction by coupling it to a highly exergonic reaction. For example, in many organisms, the synthesis of glucose 6-phosphate is the first step in the utilization of glucose. In principle, the synthesis could be accomplished by this reaction: But the positive value of ΔG′° predicts that under standard conditions the reaction will tend not to proceed spontaneously in the direction written. Another cellular reaction, the hydrolysis of ATP to ADP and Pi, is highly exergonic: G lucose+ Pi→ glucose 6-phosphate + H2O ΔG′°= 13.8KJ /mo AT P + H2O → AD P + Pi ΔG′°= −30.5 kJ /mol These two reactions share the common intermediates Pi and H2O and may be expressed as sequential reactions: The overall standard free-energy change is obtained by adding the ΔG′° values for individual reactions: ΔG′°Sum = 13.8 kJ /mol+ (−30.5 kJ /mol)= −16.7 kJ /mol The overall reaction is exergonic. In this case, energy stored in ATP is used to drive the synthesis of glucose 6- phosphate, even though its formation from glucose and inorganic phosphate (Pi) is endergonic. The pathway of glucose 6- phosphate formation from glucose by phosphoryl transfer from ATP is different from reactions (1) and (2), but the net result is the same as the sum of the two reactions. The standard free-energy change is a state function. In thermodynamic calculations, all that matters is the state of the system at the beginning of the process and its state at the end; the route between the initial and final states is immaterial. We have said that ΔG′° is a way of expressing the equilibrium constant for a reaction. For reaction (1), K′eq1= = 3.9× 10−3 M −1 Notice that H2O is not included in this expression, as its concentration (55.5 M) is assumed to remain unchanged by the reaction. The equilibrium constant for the hydrolysis of ATP is K′eq2 = = 2.0× 105M The equilibrium constant for the two coupled reactions is K′eq3 = = (K′eq1)(K′eq2)= (3.9× 10−3M −1)(2.0× 105M ) = 7.8× 102 This calculation illustrates an important point about equilibrium constants: although the ΔG′° values for two reactions that sum to a third, overall reaction are additive, the K′eq for the overall reaction is the product of the individual K′eq values for the two [glucose 6-phosphate]eq [glucose]eq[Pi]eq [AD P]eq[Pi]eq [AT P]eq [glucose6-phosphate]eq[AD P]eq[Pi]eq [glucose]eq[Pi]eq[AT P]eq reactions. Equilibrium constants are multiplicative. By coupling ATP hydrolysis to glucose 6-phosphate synthesis, the K′eq for formation of glucose 6-phosphate from glucose has been raised by a factor of about 2× 105 compared with the direct reaction between glucose and Pi. This strategy of coupling endergonic processes to exergonic reactions that drive them is employed by all living cells in the synthesis of metabolic intermediates and cellular components. Obviously, the strategy works only if compounds such as ATP are continuously available. In the following chapters we consider several of the most important cellular pathways for producing ATP. For more practice in dealing with free-energy changes and equilibrium constants for coupled reactions, see Worked Examples 1-1, 1-2, and 1-3 in Chapter 1 (pp. 24–25). SUMMARY 13.1 Bioenergetics and Thermodynamics Bioenergetics is the quantitative study of energy relationships and energy conversions in biological systems. Biological energy transformations obey the laws of thermodynamics. Living cells constantly perform work. They require energy for maintaining their highly organized structures, synthesizing cellular components, transporting small molecules and ions across membranes, and generating electric currents. All chemical reactions are influenced by two forces: the tendency to achieve the most stable bonding state (for which enthalpy, H, is a useful expression) and the tendency to achieve the highest degree of randomness, expressed as entropy, S. The driving force in a reaction is ΔG, the free-energy change, which represents the net effect of these two factors: ΔG= ΔH−TΔS. The standard transformed free-energy change, ΔG′°, is a physical constant that is characteristic for a given reaction and can be calculated from the equilibrium constant for the reaction: ΔG′°= −RT ln K′eq. The actual free-energy change, ΔG, is a variable that depends on ΔG′° and on the concentrations of reactants and products: ΔG= ΔG′°+RT ln([products]/[reactants]). When ΔG is large and negative, the reaction tends to go in the forward direction; when ΔG is large and positive, the reaction tends to go in the reverse direction; and when ΔG= 0, the system is at equilibrium. The free-energy change for a reaction is independent of the pathway by which the reaction occurs. Free-energy changes are additive; the net chemical reaction that results from successive reactions sharing a common intermediate has an overall free- energy change that is the sum of the ΔG values for the individual reactions. 13.2 Chemical Logic and Common Biochemical Reactions The biological energy transductions we are concerned with in this book are chemical reactions. Cellular chemistry does not encompass every kind of reaction learned in a typical organic chemistry course. Which reactions take place in biological systems and which do not is determined by (1) their relevance to that particular metabolic system and (2) their rates. Both considerations play major roles in shaping the metabolic pathways we consider throughout the rest of the book. A relevant reaction is one that makes use of an available substrate and converts it to a useful product. However, even a potentially relevant reaction may not occur. Some chemical transformations are too slow (have activation energies that are too high) to contribute to living systems, even with the aid of powerful enzyme catalysts. The reactions that do occur in cells represent a toolbox that evolution has used to construct metabolic pathways that circumvent the “impossible” reactions. Learning to recognize the plausible reactions can be a great aid in developing a command of biochemistry. Even so, the number of metabolic transformations taking place in a typical cell can seem overwhelming. Most cells have the capacity to carry out thousands of specific, enzyme-catalyzed reactions: for example, transformation of a simple nutrient such as glucose into amino acids, nucleotides, or lipids; extraction of energy from fuels by oxidation; and polymerization of monomeric subunits into macromolecules. Biochemical Reactions Occur in Repeating Patterns To study these reactions, some organization is essential. There are patterns within the chemistry of life; you do not need to learn every individual reaction to comprehend the molecular logic of biochemistry. Most of the reactions in living cells fall into one of five general categories: (1) reactions that make or break carbon–carbon bonds; (2) internal rearrangements, isomerizations, and eliminations; (3) free-radical reactions; (4) group transfers; and (5) oxidation-reductions. We discuss each of these in more detail below and refer to some examples of each type in later chapters. Note that the five reaction types are not mutually exclusive; for example, an isomerization reaction may involve a free-radical intermediate. Before proceeding, however, we should review two basic chemical principles. First, a covalent bond consists of a shared pair of electrons, and the bond can be broken in two general ways (Fig. 13-1). In homolytic cleavage, each atom leaves the bond as a radical, carrying one unpaired electron. In heterolytic cleavage, which is more common, one atom retains both bonding electrons. The species most o en generated when C— C and C— H bonds are cleaved are illustrated in Figure 13-1. Carbanions, carbocations, and hydride ions are highly unstable; this instability shapes the chemistry of these ions, as we shall see. FIGURE 13-1 Two mechanisms for cleavage of a C— C or C— H bond. In a homolytic cleavage, each atom keeps one of the bonding electrons, resulting in the formation of carbon radicals (carbons having unpaired electrons) or uncharged hydrogen atoms. In a heterolytic cleavage, one of the atoms retains both bonding electrons. This can result in the formation of carbanions, carbocations, protons, or hydride ions. The second basic principle is that many biochemical reactions involve interactions between nucleophiles (functional groups rich in and capable of donating electrons) and electrophiles (electron-deficient functional groups that seek electrons). Nucleophiles combine with and give up electrons to electrophiles. Common biological nucleophiles and electrophiles are shown in Figure 13-2. Note that a carbon atom can act as either a nucleophile or an electrophile, depending on which bonds and functional groups surround it. FIGURE 13-2 Common nucleophiles and electrophiles in biochemical reactions. Chemical reaction mechanisms, which trace the formation and breakage of covalent bonds, are communicated with dots and curved arrows, a convention known informally as “electron pushing.” A covalent bond consists of a shared pair of electrons. Nonbonded electrons important to the reaction mechanism are designated by dots (:). Curved arrows ( ) represent the movement of electron pairs. For movement of a single electron (as in a free-radical reaction), a single-headed (fishhook-type) arrow is used ( ). Most reaction steps involve an unshared electron pair. Reactions That Make or Break Carbon– Carbon Bonds Heterolytic cleavage of a C— C bond yields a carbanion and a carbocation (Fig. 13-1). Conversely, the formation of a C— C bond involves the combination of a nucleophilic carbanion and an electrophilic carbocation. Carbanions and carbocations are generally so unstable that their formation as reaction intermediates can be energetically unfeasible, even with enzyme catalysts. For the purpose of cellular biochemistry, they are impossible reactions — unless chemical assistance is provided in the form of functional groups containing electronegative atoms (O and N) that can alter the electronic structure of adjacent carbon atoms so as to stabilize and facilitate the formation of carbanion and carbocation intermediates. Carbonyl groups are particularly important in the chemical transformations of metabolic pathways. The carbon of a carbonyl group has a partial positive charge due to the electron- withdrawing property of the carbonyl oxygen, and so is an electrophilic carbon (Fig. 13-3a). A carbonyl group can thus facilitate the formation of a carbanion on an adjoining carbon by delocalizing the carbanion’s negative charge (Fig. 13-3b). An imine group (see Fig. 1-14) can serve a similar function (Fig. 13- 3c). The capacity of carbonyl and imine groups to delocalize electrons can be further enhanced by a general acid catalyst or by a metal ion (M e2+) such as M g2+ (Fig. 13-3d). FIGURE 13-3 Chemical properties of carbonyl groups. (a) The carbon atom of a carbonyl group is an electrophile by virtue of the electron-withdrawing capacity of the electronegative oxygen atom, which results in a structure in which the carbon has a partial positive charge. (b) Within a molecule, delocalization of electrons into a carbonyl group stabilizes a carbanion on an adjacent carbon, facilitating its formation. (c) Imines function much like carbonyl groups in facilitating electron withdrawal. (d) Carbonyl groups do not always function alone; their capacity as electron sinks o en is augmented by interaction with either a metal ion (M e2+, such as M g2+) or a general acid (HA). The importance of a carbonyl group is evident in three major classes of reactions in which C— C bonds are formed or broken (Fig. 13-4): aldol condensations, Claisen ester condensations, and decarboxylations. In each type of reaction, a carbanion intermediate is stabilized by a carbonyl group, and in many cases another carbonyl provides the electrophile with which the nucleophilic carbanion reacts. FIGURE 13-4 Some common reactions that form and break C— C bonds in biological systems. For both the aldol condensation and the Claisen condensation, a carbanion serves as nucleophile and the carbon of a carbonyl group serves as electrophile. The carbanion is stabilized in each case by another carbonyl at the adjoining carbon. In the decarboxylation reaction, a carbanion is formed on the carbon shaded blue as the CO2 leaves. The reaction would not occur at an appreciable rate without the stabilizing effect of the carbonyl adjacent to the carbanion carbon. Wherever a carbanion is shown, a stabilizing resonance with the adjacent carbonyl, as shown in Figure 13-3b, is assumed. An imine (Fig. 13-3c) or other electron-withdrawing group (including certain enzymatic cofactors such as pyridoxal) can replace the carbonyl group in the stabilization of carbanions. An aldol condensation is a common route to the formation of a C— C bond; the aldolase reaction, which converts a six-carbon compound to two three-carbon compounds in glycolysis, is an aldol condensation in reverse (see Fig. 14-5). In a Claisen condensation, the carbanion is stabilized by the carbonyl of an adjacent thioester; an example is the synthesis of citrate in the citric acid cycle (see Fig. 16-9). Decarboxylation also commonly involves the formation of a carbanion stabilized by a carbonyl group; the acetoacetate decarboxylase reaction that occurs in the formation of ketone bodies during fatty acid catabolism provides an example (see Fig. 17-16). Entire metabolic pathways are organized around the introduction of a carbonyl group in a particular location so that a nearby carbon–carbon bond can be formed or cleaved. In some reactions, an imine or a specialized cofactor such as pyridoxal phosphate plays the electron- withdrawing role, instead of a carbonyl group. The carbocation intermediate occurring in some reactions that form or cleave C— C bonds is generated by the elimination of an excellent leaving group, such as pyrophosphate (see “Group Transfer Reactions” below). An example is the prenyltransferase reaction (Fig. 13-5), an early step in the pathway of cholesterol biosynthesis. FIGURE 13-5 Carbocations in carbon–carbon bond formation. In one of the early steps in cholesterol biosynthesis, the enzyme prenyltransferase catalyzes condensation of isopentenyl pyrophosphate and dimethylallyl pyrophosphate to form geranyl pyrophosphate (see Fig. 21-36). The reaction is initiated by elimination of pyrophosphate from the dimethylallyl pyrophosphate to generate a carbocation, stabilized by resonance with the adjacent C═C bond. Internal Rearrangements, Isomerizations, and Eliminations Another common type of cellular reaction is an intramolecular rearrangement in which redistribution of electrons results in alterations of many different types without a change in the overall oxidation state of the molecule. For example, different groups in a molecule may undergo oxidation-reduction, with no net change in oxidation state of the molecule; groups at a double bond may undergo a cis-trans rearrangement; or the positions of double bonds may be transposed. An example of an isomerization entailing internal oxidation-reduction is the formation of fructose 6-phosphate from glucose 6-phosphate in glycolysis (Fig. 13-6; this reaction is discussed in detail in Chapter 14): C-1 is reduced (aldehyde to alcohol) and C-2 is oxidized (alcohol to ketone). Figure 13-6b shows the details of the electron movements in this type of isomerization. A cis-trans rearrangement is illustrated by the prolyl cis-trans isomerase reaction in the folding of certain proteins (see p. 133). A simple transposition of a C═C bond occurs during metabolism of oleic acid, a common fatty acid (see Fig. 17-10). Some spectacular examples of double-bond repositioning occur in the biosynthesis of cholesterol (see Fig. 21- 37). FIGURE 13-6 Isomerization and elimination reactions. (a) The conversion of glucose 6- phosphate to fructose 6-phosphate, a reaction of sugar metabolism catalyzed by phosphohexose isomerase. (b) This reaction proceeds through an enediol intermediate. Light red screens follow the path of oxidation from le to right. B1 and B2 are ionizable groups on the enzyme; they are capable of donating and accepting protons (acting as general acids or general bases) as the reaction proceeds. An example of an elimination reaction that does not affect overall oxidation state is the loss of water from an alcohol, resulting in the introduction of a C═C bond: Similar reactions can result from eliminations in amines. Free-Radical Reactions Once thought to be rare, the homolytic cleavage of covalent bonds to generate free radicals has now been found in a wide range of biochemical processes. These include isomerizations that make use of adenosylcobalamin (vitamin B12) or S- adenosylmethionine, which are initiated with a 5′-deoxyadenosyl radical (see the methylmalonyl-CoA mutase reaction in Box 17-2); certain radical-initiated decarboxylation reactions (Fig. 13-7); some reductase reactions, such as that catalyzed by ribonucleotide reductase (see Fig. 22-43); and some rearrangement reactions, such as that catalyzed by DNA photolyase (see Fig. 25-25). FIGURE 13-7 A free radical–initiated decarboxylation reaction. The biosynthesis of heme in Escherichia coli includes a decarboxylation step in which propionyl side chains on the coproporphyrinogen III intermediate are converted to the vinyl side chains of protoporphyrinogen IX. When the bacteria are grown anaerobically the enzyme oxygen-independent coproporphyrinogen III oxidase, also called HemN protein, promotes decarboxylation via the free-radical mechanism shown here. The acceptor of the released electron is not known. For simplicity, only the relevant portions of the large coproporphyrinogen III and protoporphyrinogen molecules are shown; the entire structures are given in Figure 22-26. When E. coli is grown in the presence of oxygen, this reaction is an oxidative decarboxylation and is catalyzed by a different enzyme. [Information from G. Layer et al., Curr. Opin. Chem. Biol. 8:468, 2004, Fig. 4.] Group Transfer Reactions The transfer of acyl, glycosyl, and phosphoryl groups from one nucleophile to another is common in living cells. Acyl group transfer generally involves the addition of a nucleophile to the carbonyl carbon of an acyl group to form a tetrahedral intermediate: The chymotrypsin reaction is one example of acyl group transfer (see Fig. 6-27). Glycosyl group transfers involve nucleophilic substitution at C-1 of a sugar ring, which is the central atom of an acetal. In principle, the substitution could proceed by an SN1 or SN2 pathway. Phosphoryl group transfers play a special role in metabolic pathways, and these transfer reactions are discussed in detail in Section 13.3. A general theme in metabolism is the attachment of a good leaving group to a metabolic intermediate to “activate” the intermediate for subsequent reaction. Among the better leaving groups in nucleophilic substitution reactions are inorganic orthophosphate (the ionized form of H3PO4 at neutral pH, a mixture of H2PO− 4 and HPO2− 4 , commonly abbreviated Pi) and inorganic pyrophosphate (P2O4− 7 , abbreviated PPi); esters and anhydrides of phosphoric acid are effectively activated for reaction. Nucleophilic substitution is made more favorable by the attachment of a phosphoryl group to an otherwise poor leaving group such as — OH. Nucleophilic substitutions in which the phosphoryl group (— PO2− 3 ) serves as a leaving group occur in hundreds of metabolic reactions. Phosphorus can form five covalent bonds. The conventional representation of Pi (Fig. 13-8a), with three P—O bonds and one P═O bond, is a convenient but inaccurate picture. In Pi, four equivalent phosphorus–oxygen bonds share some double-bond character, and the anion has a tetrahedral structure (Fig. 13-8b). Because oxygen is more electronegative than phosphorus, the sharing of electrons is unequal: the central phosphorus bears a partial positive charge and can therefore act as an electrophile. In a great many metabolic reactions, a phosphoryl group (— PO2− 3 ) is transferred from ATP to an alcohol, forming a phosphate ester (Fig. 13-8c), or to a carboxylic acid, forming a mixed anhydride. When a nucleophile attacks the electrophilic phosphorus atom in ATP, a relatively stable pentacovalent structure forms as a reaction intermediate (Fig. 13-8d). With departure of the leaving group (ADP), the transfer of a phosphoryl group is complete. The large family of enzymes that catalyze phosphoryl group transfers with ATP as donor are called kinases (Greek kinein, “to move”). Hexokinase, for example, “moves” a phosphoryl group from ATP to the hexose glucose. Box 13-1 offers a primer on some of the broad classes of enzymes (including kinases) that you will encounter in your study of metabolism.
FIGURE 13-8 Phosphoryl group transfers: some of the participants. (a) In one (inadequate) representation of Pi, three oxygens are single-bonded to phosphorus, and the fourth is double-bonded, allowing the four different resonance structures shown here. (b) The resonance structures of Pi can be represented more accurately by showing all four phosphorus–oxygen bonds with some double-bond character; the hybrid orbitals so represented are arranged in a tetrahedron with P at its center. (c) When a nucleophile Z (in this case, the — OH on C-6 of glucose) attacks ATP, it displaces ADP (W). In this SN2 reaction, a pentacovalent intermediate (d) forms transiently. BOX 13-1 A Primer on Enzyme Names The name kinase is applied to enzymes that transfer a phosphoryl group from a nucleoside triphosphate such as ATP to an acceptor molecule—a sugar (as in hexokinase and glucokinase), a protein (as in glycogen phosphorylase kinase), another nucleotide (as in nucleoside diphosphate kinase), or a metabolic intermediate such as oxaloacetate (as in PEP carboxykinase). The reaction catalyzed by a kinase is a phosphorylation. On the other hand, phosphorolysis is a displacement reaction in which phosphate is the attacking species and becomes covalently attached at the point of bond breakage. Such reactions are catalyzed by phosphorylases. Glycogen phosphorylase, for example, catalyzes the phosphorolysis of glycogen, producing glucose 1-phosphate. Dephosphorylation, the removal of a phosphoryl group from a phosphate ester, is catalyzed by phosphatases, with water as the attacking species. Fructose bisphosphatase-1 converts fructose 1,6-bisphosphate to fructose 6- phosphate in gluconeogenesis, and phosphorylase a phosphatase removes phosphoryl groups from phosphoserine in phosphorylated glycogen phosphorylase. Whew! Citrate synthase, the first enzyme in the citric acid cycle (see Fig. 16-7), is one of many enzymes that catalyze condensation reactions, yielding a product more chemically complex than its precursors. Synthases catalyze condensation reactions in which no nucleoside triphosphate (ATP, GTP, and so forth) is required as an energy source. Synthetases catalyze condensations that do use ATP or another nucleoside triphosphate as a source of energy for the synthetic reaction. Succinyl-CoA synthetase is such an enzyme. Ligases (from the Latin ligare, “to tie together”) are enzymes that catalyze condensation reactions in which two atoms are joined, using ATP or another energy source. (Thus, synthetases are ligases.) DNA ligase, for example, closes breaks in DNA molecules, using energy supplied by either ATP or NAD+; it is widely used in joining DNA pieces for genetic engineering. Ligases are not to be confused with lyases, enzymes that catalyze cleavages (or, in the reverse direction, additions) in which electronic rearrangements occur. The PDH complex, which oxidatively cleaves CO2 from pyruvate, is a member of the large class of lyases. In some biological oxidation reactions, molecular oxygen is the electron acceptor. If the oxygen atoms do not appear in the oxidized product, the enzyme is an oxidase. If one or both of the oxygen atoms do appear in the oxidized product, as a new hydroxyl or carboxyl group, for example, the enzyme is an oxygenase. Both classes are further subdivided. Mixed-function oxidases oxidize two different substrates simultaneously. Monooxygenases and dioxygenases catalyze reactions in which one or two oxygen atoms, respectively, are incorporated into the organic product. These enzymes are particularly important in biosynthetic pathways of fatty acids and eicosanoids (see Box 21-1). Dehydrogenases catalyze oxidation-reductase reactions in which NAD+ is electron acceptor, and molecular oxygen is generally not involved. Unfortunately, these descriptions of enzyme types overlap, and many enzymes have two or more common names. Succinyl-CoA synthetase, for example, is also called succinate thiokinase; the enzyme is both a synthetase in the citric acid cycle and a kinase when acting in the direction of succinyl-CoA synthesis. This raises another source of confusion in the naming of enzymes. An enzyme may have been discovered by the use of an assay in which, say, A is converted to B. The enzyme is then named for that reaction. Later work may show, however, that in the cell, the enzyme functions primarily in converting B to A. Commonly, the first name continues to be used, although the metabolic role of the enzyme would be better described by naming it for the reverse reaction. The glycolytic enzyme pyruvate kinase illustrates this situation (p. 521). To a beginner in biochemistry, this duplication in nomenclature can be bewildering. International committees have made heroic efforts to systematize the nomenclature of enzymes (see Table 6-3 for a brief summary of the system), but some systematic names have proved too long and cumbersome and are not frequently used in biochemical conversation. We have tried throughout this book to use the enzyme name most commonly employed by working biochemists and to point out cases in which an enzyme has more than one widely used name. Phosphoryl groups are not the only groups that activate molecules for reaction. Thioalcohols (thiols), in which the oxygen atom of an alcohol is replaced with a sulfur atom, are also good leaving groups. Thiols activate carboxylic acids by forming thioesters (thiol esters). In later chapters we discuss several reactions, including those catalyzed by the fatty acyl synthases in lipid synthesis (see Fig. 21-2), in which nucleophilic substitution at the carbonyl carbon of a thioester results in transfer of the acyl group to another moiety. Oxidation-Reduction Reactions We will encounter carbon atoms in five oxidation states, depending on the elements with which they share electrons (Fig. 13-9), and transitions between these states are of crucial importance in metabolism (oxidation-reduction reactions are the topic of Section 13.4). In many biological oxidations, a compound loses two electrons and two hydrogen ions (that is, two hydrogen atoms); these reactions are commonly called dehydrogenations, and the enzymes that catalyze them are called dehydrogenases (Fig. 13-10). In some, but not all, biological oxidations, a carbon atom becomes covalently bonded to an oxygen atom. The enzymes that catalyze oxidations with oxygen as electron acceptor are generally called oxidases or, if the oxygen atom is derived directly from molecular oxygen (O2) and incorporated into the product, oxygenases. FIGURE 13-9 The oxidation levels of carbon in biomolecules. Each compound is formed by oxidation of the carbon shown in red in the compound immediately above. Carbon dioxide is the most highly oxidized form of carbon found in living systems. FIGURE 13-10 An oxidation-reduction reaction. Shown here is the oxidation of lactate to pyruvate. In this dehydrogenation, two electrons and two hydrogen ions (the equivalent of two hydrogen atoms) are removed from C-2 of lactate, an alcohol, to form pyruvate, a ketone. In cells the reaction is catalyzed by lactate dehydrogenase and the electrons are transferred to the cofactor nicotinamide adenine dinucleotide (NAD+). This reaction is fully reversible; pyruvate can be reduced by electrons transferred from the cofactor. Every oxidation must be accompanied by a reduction, in which an electron acceptor acquires the electrons removed by oxidation. Oxidation reactions generally release energy (think of campfires: the compounds in wood are oxidized by oxygen molecules in the air). Most living cells obtain the energy needed for cellular work by oxidizing metabolic fuels such as carbohydrates or fat (photosynthetic organisms can also trap and use the energy of sunlight). The catabolic (energy-yielding) pathways described in Chapters 14 through 19 are oxidative reaction sequences that result in the transfer of electrons from fuel molecules, through a series of electron carriers, to oxygen. The high affinity of O2 for electrons makes the overall electron-transfer process highly exergonic, providing the energy that drives ATP synthesis — the central goal of catabolism. Many of the reactions within these five classes are facilitated by cofactors, in the form of coenzymes and metal ions (vitamin B12, S-adenosylmethionine, folate, nicotinamide, and Fe2+ are some examples). Cofactors bind to enzymes — in some cases reversibly, in other cases almost irreversibly — and give them the capacity to promote a particular kind of chemistry (p. 178). Most cofactors participate in a narrow range of closely related reactions. In the following chapters, we introduce and discuss each important cofactor at the point where we first encounter its function. The cofactors provide another way to organize the study of biochemical processes, given that the reactions facilitated by a given cofactor generally are mechanistically related. Biochemical and Chemical Equations Are Not Identical Biochemists write metabolic equations in a simplified way, and this is particularly evident for reactions involving ATP. Phosphorylated compounds can exist in several ionization states and, as we have noted, the different species can bind M g2+. For example, at pH 7 and 2 mM M g2+, ATP exists as an equilibrium distribution of the forms ATP4−, HATP3−, H2ATP2−, M gHATP−, and M g2ATP. In thinking about the biological role of ATP, however, we are not always interested in all this detail, and so we consider ATP as an entity made up of a sum of species, and we write its hydrolysis as the biochemical equation ATP+ H2O → ADP+ Pi where ATP, ADP, and Pi are sums of species. The corresponding standard transformed equilibrium constant, K′eq = [ADP]eq[Pi]eq/[ATP]eq, depends on the pH and the concentration of free M g2+. Note that H+ and M g2+ do not appear in the biochemical equation, because their concentrations are not significantly changed by the reaction. Thus a biochemical equation does not necessarily balance H, Mg, or charge, although it does balance all other elements involved in the reaction (C, N, O, and P in the equation above). We can write a chemical equation that does balance for all elements and for charge. For example, when ATP is hydrolyzed at a pH above 8.5 in the absence of M g2+, the chemical reaction is represented by ATP4− + H2O → ADP3− + HPO2− 4 + H+ The corresponding equilibrium constant, Keq = [ADP3−]eq[HPO2− 4 ]eq[H+]eq/[ATP4−]eq, depends only on temperature, pressure, and ionic strength. Both ways of writing a metabolic reaction have value in biochemistry. Chemical equations are needed when we want to account for all atoms and charges in a reaction, as when we are considering the mechanism of a chemical reaction. Biochemical equations are used to determine in which direction a reaction will proceed spontaneously, given a specified pH and [M g2+], or to calculate the equilibrium constant of such a reaction. Throughout this book we use biochemical equations, unless the focus is on chemical mechanism, and we use values of ΔG′° and K′eqas determined at pH 7 and 1 mMM g2+. SUMMARY 13.2 Chemical Logic and Common Biochemical Reactions Living systems make use of a large number of chemical reactions that can be classified into five general types: reactions that make or break carbon–carbon bonds; internal rearrangements and eliminations; free-radical reactions; group transfers; and oxidation-reduction reactions. Heterolytic cleavages occur o en in reactions that make or break C— C bonds. Carbonyl groups play a special role in reactions that form or cleave C— C bonds. Carbanion intermediates are common and are stabilized by adjacent carbonyl groups or, less o en, by imines or certain cofactors. A redistribution of electrons can produce internal rearrangements, isomerizations, and eliminations. Such reactions include intramolecular oxidation-reduction, change in cis-trans arrangement at a double bond, and transposition of double bonds. Homolytic cleavage of covalent bonds to generate free radicals occurs in some pathways. Phosphoryl transfer reactions are an especially important type of group transfer in cells, required for the activation of molecules for reactions that would otherwise be highly unfavorable. Oxidation-reduction reactions involve the loss or gain of electrons: one reactant gains electrons and is reduced, while the other loses electrons and is oxidized. Oxidation reactions generally release energy and are important in catabolism. Biochemists o en write reaction equations that are not balanced for H+ and don’t attempt to describe the state of phosphate ionization. 13.3 Phosphoryl Group Transfers and ATP Having developed some fundamental principles of energy changes in chemical systems and reviewed the common classes of reactions, we can now examine the energy cycle in cells and the special role of ATP as the energy currency that links catabolism and anabolism (see Fig. 1-28). Heterotrophic cells obtain free energy in a chemical form by the catabolism of nutrient molecules, and they use that energy to make ATP from ADP and Pi. ATP then donates some of its chemical energy to endergonic processes such as the synthesis of metabolic intermediates and macromolecules from smaller precursors, the transport of substances across membranes against concentration gradients, and mechanical motion. This donation of energy from ATP generally involves the covalent participation of ATP in the reaction that is to be driven, with the eventual result that ATP is converted to ADP and Pi or, in some reactions, to AMP and 2 Pi. We discuss here the chemical basis for the large free-energy changes that accompany hydrolysis of ATP and other high-energy phosphate compounds, and we show that most cases of energy donation by ATP involve group transfer, not simple hydrolysis of ATP. To illustrate the range of energy transductions in which ATP provides the energy, we consider the synthesis of information- rich macromolecules, the transport of solutes across membranes, and motion produced by muscle contraction. The Free-Energy Change for ATP Hydrolysis Is Large and Negative Figure 13-11 summarizes the chemical basis for the relatively large, negative, standard free energy of hydrolysis of ATP. The hydrolytic cleavage of the terminal phosphoric acid anhydride (phosphoanhydride) bond in ATP separates one of the three negatively charged phosphates and thus relieves some of the internal electrostatic repulsion in ATP; the Pi released is stabilized by the formation of several resonance forms not possible in ATP.
FIGURE 13-11 Chemical basis for the large free-energy change associated with ATP hydrolysis. The charge separation that results from hydrolysis relieves electrostatic repulsion among the four negative charges on ATP. The product inorganic phosphate (Pi) is stabilized by formation of a resonance hybrid, in which each of the four phosphorus–oxygen bonds has the same degree of double-bond character and the hydrogen ion is not permanently associated with any one of the oxygens. (Some degree of resonance stabilization also occurs in phosphates involved in ester or anhydride linkages, but fewer resonance forms are possible than for Pi.) A third factor (not shown) that favors ATP hydrolysis is the greater degree of solvation (hydration) of the products Pi and ADP relative to ATP, which further stabilizes the products relative to the reactants. The free-energy change for ATP hydrolysis is −30.5 kJ /mol under standard conditions, but the actual free energy of hydrolysis (ΔG) of ATP in living cells is very different: the cellular concentrations of ATP, ADP, and Pi are not identical and are much lower than the 1.0 M of standard conditions (Table 13-5). Furthermore, M g2+ in the cytosol binds to ATP and ADP (Fig. 13-12), and for most enzymatic reactions that involve ATP as phosphoryl group donor, the true substrate is M gAT P2−. The relevant ΔG′° is therefore that for M gAT P2− hydrolysis. We can calculate ΔG for ATP hydrolysis using data such as those in Table 13-5. The actual free energy of hydrolysis of ATP under intracellular conditions is o en called its phosphorylation potential, ΔGp, for reasons we will explain. TABLE 13-5 Total Concentrations of Adenine Nucleotides, Inorganic Phosphate, and Phosphocreatine in Some Cells Concentration (m ) Cell type ATP ADP AMP Pi PCr Rat hepatocyte 3.38 1.32 0.29 4.8 0 Rat myocyte 8.05 0.93 0.04 8.05 27 Rat neuron 2.59 0.73 0.06 2.72 4.7 Human erythrocyte 2.25 0.25 0.02 1.65 0 E. coli cell 9.6 0.56 0.28 — — For erythrocytes, the concentrations are those of the cytosol (human erythrocytes lack a nucleus and mitochondria). In the other types of cells, the data are for the entire cell contents, although the cytosol and the mitochondria have very different concentrations of ADP. PCr is phosphocreatine, discussed on p. 487. This value reflects total concentration; the true value for free ADP may be much lower (Worked Example 13-2). Mammalian data from R. L. Veech et al., J. Biol. Chem. 254:6538, 1979. E. coli data from B. D. Bennett et al., Nat. Chem. Biol. 5:593, 2009. a b a b FIGURE 13-12 M g2+ and ATP. Formation of M g2+ complexes partially shields the negative charges and influences the conformation of the phosphate groups in nucleotides such as ATP and ADP. Because the concentrations of ATP, ADP, and Pi differ from one cell type to another, ΔGp for ATP likewise differs among cells. Moreover, in any given cell, ΔGp can vary from time to time, depending on the metabolic conditions and how they influence the concentrations of ATP, ADP, Pi, and H+ (pH). We can calculate the actual free-energy change for any given metabolic reaction as it occurs in a cell, provided we know the concentrations of all the reactants and products and other factors (such as pH, temperature, and [M g2+]) that may affect the actual free-energy change. WORKED EXAMPLE 13-2 Calculation of ΔGp Calculate the actual free energy of hydrolysis of ATP, ΔGp, in human erythrocytes. The standard free energy of hydrolysis of ATP is −30.5 kJ /mol, and the concentrations of ATP, ADP, and Pi in erythrocytes are as shown in Table 13-5. Assume that the pH is 7.0 and the temperature is 37 °C (human body temperature). What does this reveal about the amount of energy required to synthesize ATP under the same cellular conditions? SOLUTION: The concentrations of ATP, ADP, and Pi in human erythrocytes are 2.25, 0.25, and 1.65 mM, respectively. The actual free energy of hydrolysis of ATP under these conditions is given by the relationship (see Eqn 13-4) ΔGp = ΔG′°+ RT ln Substituting the appropriate values, we get [AD P][Pi] [AT P] Thus ΔGp, the actual free-energy change for ATP hydrolysis in the intact erythrocyte (−52 kJ /mol), is much larger than the standard free-energy change (−30.5 kJ /mol). Similarly, the free energy required to synthesize ATP from ADP and Pi under the conditions prevailing in the erythrocyte would be 52 kJ/mol. To further complicate the issue, the total concentrations of ATP, ADP, and Pi (and H+) in a cell — such as the values given in Table 13-5 — may be substantially higher than the free concentrations, which are the thermodynamically relevant values. The difference is due to tight binding of ATP, ADP, and Pi to cellular proteins. For example, the free [ADP] in resting muscle has been variously estimated at between 1 and 37 µM. Using the value 25 µM in Worked Example 13-2, we would get a ΔGp of −58 kJ /mol. Calculation of the exact value of ΔGp, however, is perhaps less instructive than the generalization we can make about actual free-energy changes: in vivo, the energy released by ATP hydrolysis is greater than the standard free-energy change, ΔG′°. ΔGp = −30.5 kJ /mol+ [(8.315 J /mol∙K)(310 K) ln = −30.5 kJ /mol+ (2.58 kJ /mol) ln 1.8× 10−4 = −30.5 kJ /mol+ (2.58 kJ /mol)(−8.6) = −30.5 kJ /mol− 22 kJ /mol = −52 kJ /mol (0.25× 10 (2.2 In the following discussions we use the ΔG′° value for ATP hydrolysis because this allows comparisons, on the same basis, with the energetics of other cellular reactions. Always keep in mind, however, that in living cells ΔG is the relevant quantity — for ATP hydrolysis and all other reactions — and may be quite different from ΔG′°. Here we must make an important point about cellular ATP levels. We have shown (and will discuss further) how the chemical properties of ATP make it a suitable form of energy currency in cells. But it is not merely the molecule’s intrinsic chemical properties that give it this ability to drive metabolic reactions and other energy-requiring processes. Even more important is that, in the course of evolution, there has been a very strong selective pressure for regulatory mechanisms that hold cellular ATP concentrations far above the equilibrium concentrations for the hydrolysis reaction. When the ATP level drops, not only does the amount of fuel decrease, but the fuel itself loses its potency: ΔG for its hydrolysis (that is, its phosphorylation potential, ΔGp) is diminished. As our discussions of the metabolic pathways that produce and consume ATP will show, living cells have developed elaborate mechanisms — o en at what might seem to us the expense of efficiency — to maintain high concentrations of ATP. Other Phosphorylated Compounds and Thioesters Also Have Large, Negative Free Energies of Hydrolysis ATP is not the only biological compound with a large, negative free energy of hydrolysis. Table 13-6 lists the standard free energies of hydrolysis for some biologically important phosphorylated compounds. In all of these phosphate-releasing reactions, a few of which we describe below, the several resonance forms available to Pi (Fig. 13-11) stabilize this product relative to the reactant, contributing to an already negative free- energy change. TABLE 13-6 Standard Free Energies of Hydrolysis of Some Phosphorylated Compounds and Acetyl-CoA (a Thioester) ΔG′° Compounds and Acetyl-CoA (kJ/mol) (kcal/mol) Phosphoenolpyruvate −61.9 −14.8 1,3-Bisphosphoglycerate (→ 3-phosphoglycerate+ Pi) −49.3 −11.8 Phosphocreatine −43.0 −10.3 ADP (→ AM P + Pi) −32.8 −7.8 ATP (→ AD P + Pi) −30.5 −7.3 ATP (→ AM P + PPi) −45.6 −10.9 AMP (→ adenosine+ Pi) −14.2 −3.4 PPi(→ 2Pi) −19.2 −4.0 Glucose 3-phosphate −20.9 −5.0 Fructose 6-phosphate −15.9 −3.8 Glucose 6-phosphate −13.8 −3.3 Glycerol 3-phosphate −9.2 −2.2 Acetyl-CoA −31.4 −7.5 Data mostly from W. P. Jencks, in Handbook of Biochemistry and Molecular Biology, 3rd edn (G. D. Fasman, ed.), Physical and Chemical Data, Vol. 1, p. 296, CRC Press, 1976. Value for the free energy of hydrolysis of PPi from P. A. Frey and A. Arabshahi, Biochemistry 34:11,307, 1995. Phosphoenolpyruvate (PEP; Fig. 13-13), a central intermediate in the energy-conserving process of glycolysis (Chapter 14), contains a phosphate ester bond that undergoes hydrolysis to yield the enol form of pyruvate, and this direct product can tautomerize to the more stable keto form. Because the reactant (PEP) has only one form (enol) and the product (pyruvate) has two possible forms, the reaction occurs with a gain in entropy, and the product is therefore stabilized relative to the reactant. This is a major contributor to the high standard free energy of hydrolysis of phosphoenolpyruvate: ΔG′°= −61.9 kJ /mol. A second contributor is the greater resonance stabilization in the Pi released by cleavage of PEP.
FIGURE 13-13 Hydrolysis of phosphoenolpyruvate (PEP). Catalyzed by pyruvate kinase, this reaction is followed by spontaneous tautomerization of the product, pyruvate. Tautomerization is not possible in PEP, and thus the products of hydrolysis are stabilized relative to the reactants. Resonance stabilization of Pi also occurs, as shown in Figure 13-11. Another intermediate in glycolysis, the three-carbon compound 1,3-bisphosphoglycerate (Fig. 13-14), contains an anhydride bond between the C-1 carboxyl group and phosphoric acid. Hydrolysis of this acyl phosphate is accompanied by a large, negative, standard free-energy change (ΔG′°= −49.3 kJ /mol), which can, again, be explained in terms of the structure of reactant and products. When H2O is added across the anhydride bond of 1,3- bisphosphoglycerate, one of the direct products, 3- phosphoglyceric acid, can lose a proton to give the carboxylate ion, 3-phosphoglycerate, which has two equally probable resonance forms. Removal of the direct product (3- phosphoglyceric acid) by its further metabolism and formation of the resonance-stabilized ion both favor the forward reaction.
FIGURE 13-14 Hydrolysis of 1,3-bisphosphoglycerate. The direct product of hydrolysis is 3-phosphoglyceric acid, with an undissociated carboxylic acid. Its dissociation allows resonance structures that stabilize the product relative to the reactants. Resonance stabilization of Pi further contributes to the negative free-energy change. In phosphocreatine (Fig. 13-15), which is used in muscle tissue to replenish ATP a er its use in contraction, the P— N bond can be hydrolyzed to generate free creatine and Pi. The release of Pi and the resonance stabilization of creatine favor the forward reaction. The standard free-energy change of phosphocreatine hydrolysis is therefore large: −43.0 kJ /mol. FIGURE 13-15 Hydrolysis of phosphocreatine. Breakage of the P— N bond in phosphocreatine produces creatine, which is stabilized by formation of a resonance hybrid. The other product, Pi, is also resonance stabilized. Thioesters, in which a sulfur atom replaces the usual oxygen in the ester bond, also have large, negative, standard free energies of hydrolysis. Acetyl-coenzyme A, or acetyl-CoA (Fig. 13-16), is one of many thioesters important in metabolism. The acyl group in these compounds is activated for transacylation and condensation. Thioesters undergo much less resonance stabilization than do oxygen esters; consequently, the difference in free energy between the reactant and its hydrolysis products, which are resonance-stabilized, is greater for thioesters than for comparable oxygen esters (Fig. 13-17). In both cases, hydrolysis of the ester generates a carboxylic acid, which can ionize and assume several resonance forms. Together, these factors result in the large, negative ΔG′° (−31.4 kJ /mol) for acetyl-CoA hydrolysis.
FIGURE 13-16 Hydrolysis of acetyl-coenzyme A. Acetyl-CoA is a thioester with a large, negative, standard free energy of hydrolysis. Thioesters contain a sulfur atom in the position occupied by an oxygen atom in oxygen esters. The complete structure of coenzyme A (CoA, or CoASH) is shown in Figure 8-41. FIGURE 13-17 Free energy of hydrolysis for thioesters and oxygen esters. The products of both types of hydrolysis reaction have about the same free- energy content (G), but the thioester has a higher free-energy content than the oxygen ester. Orbital overlap between the O and C atoms allows resonance stabilization in oxygen esters; orbital overlap between S and C atoms is poorer and provides little resonance stabilization. To summarize, for hydrolysis reactions with large, negative, standard free-energy changes, the products are more stable than the reactants for one or more of the following reasons: (1) the bond strain in reactants due to electrostatic repulsion is relieved by charge separation, as for ATP; (2) the products are stabilized by ionization, as for ATP, acyl phosphates, and thioesters; (3) the products are stabilized by isomerization (tautomerization), as for PEP; and/or (4) the products are stabilized by resonance, as for creatine released from phosphocreatine, carboxylate ion released from acyl phosphates and thioesters, and phosphate (Pi) released from anhydride or ester linkages. ATP Provides Energy by Group Transfers, Not by Simple Hydrolysis Throughout this book you will encounter reactions or processes for which ATP supplies energy, and the contribution of ATP to these reactions is commonly indicated as in Figure 13-18a, with a single arrow showing the conversion of ATP to ADP and Pi (or, in some cases, of ATP to AMP and pyrophosphate, PPi). When written this way, these reactions of ATP seem to be simple hydrolysis reactions in which water displaces Pi (or PPi), and one is tempted to say that an ATP-dependent reaction is “driven by the hydrolysis of ATP.” This is not the case. ATP hydrolysis per se usually accomplishes nothing but the liberation of heat, which cannot drive a chemical process in an isothermal system. A single reaction arrow such as that in Figure 13-18a almost invariably represents a two-step process (Fig. 13-18b) in which part of the ATP molecule, a phosphoryl or pyrophosphoryl group or the adenylate moiety (AMP), is first transferred to a substrate molecule or to an amino acid residue in an enzyme, becoming covalently attached to the substrate or the enzyme and raising its free-energy content (activating it). Then, in a second step, the phosphate-containing moiety transferred in the first step is displaced, generating Pi, PPi, or AMP as the leaving group. Thus, ATP participates covalently in the enzyme-catalyzed reaction to which it contributes free energy. FIGURE 13-18 ATP hydrolysis in two steps. (a) The contribution of ATP to a reaction is o en shown as a single step but is almost always a two-step process. (b) Shown here is the reaction catalyzed by ATP-dependent glutamine synthetase. A phosphoryl group is transferred from ATP to glutamate, then the phosphoryl group is displaced by NH3 and released as Pi. Some processes do involve direct hydrolysis of ATP (or the related GTP, guanosine triphosphate), however. For example, noncovalent binding of ATP (or GTP), followed by its hydrolysis to ADP (or GDP, guanosine diphosphate) and Pi, can provide the energy to cycle some proteins between two conformations, producing mechanical motion. This occurs in muscle contraction (see Fig. 5-29) and in the movement of enzymes along DNA (see Fig. 25-30) or of ribosomes along messenger RNA (see Fig. 27-30). The energy-dependent reactions catalyzed by helicases, RecA protein, and some topoisomerases (Chapter 25) also involve direct hydrolysis of phosphoanhydride bonds. The AAA+ ATPases involved in DNA replication and other processes described in Chapter 25 use ATP hydrolysis to cycle associated proteins between active and inactive forms. GTP-binding proteins that act in signaling pathways directly hydrolyze GTP to drive conformational changes that terminate signals triggered by hormones or by other extracellular factors (Chapter 12). The phosphate compounds found in living organisms can be divided, somewhat arbitrarily, into two groups, based on their standard free energies of hydrolysis (Fig. 13-19). “High- energy” compounds have a ΔG′° of hydrolysis more negative than −25 kJ /mol; “low-energy” compounds have a less negative ΔG′°. Based on this criterion, ATP, with a ΔG′° of hydrolysis of −30.5 kJ /mol (−7.3 kcal/mol), is a high-energy compound; glucose 6-phosphate, with a ΔG′° of hydrolysis of −13.8 kJ /mol (−3.3 kcal/mol), is a low-energy compound. FIGURE 13-19 Ranking of biological phosphate compounds by standard free energies of hydrolysis. Phosphoryl groups, represented by , flow from high-energy phosphoryl group donors via ATP to acceptor molecules (such as glucose and glycerol) to form their low-energy phosphate derivatives. (The location of each compound’s donor phosphoryl group along the scale is an approximate indication of the compound’s ΔG′° of hydrolysis.) This flow of phosphoryl groups, catalyzed by kinases, proceeds with an overall loss of free energy under intracellular conditions. Hydrolysis of low-energy phosphate compounds releases Pi, which has an even lower phosphoryl group transfer potential. The term “high-energy phosphate bond,” long used by biochemists to describe the P— O bond broken in hydrolysis reactions, is incorrect and misleading, as it wrongly suggests that the bond itself contains the energy. In fact, the breaking of all chemical bonds requires an input of energy. The free energy released by hydrolysis of phosphate compounds does not come from the specific bond that is broken; it results from the products of the reaction having a lower free-energy content than the reactants. For simplicity, we sometimes use the term “high- energy phosphate compound” when referring to ATP or other phosphate compounds with a large, negative, standard free energy of hydrolysis. As is evident from the additivity of free-energy changes of sequential reactions (see Section 13.1), any phosphorylated compound can be synthesized by coupling the synthesis to the breakdown of another phosphorylated compound with a more negative free energy of hydrolysis. For example, because cleavage of Pi from phosphoenolpyruvate releases more energy than is needed to drive the condensation of Pi with ADP, the direct donation of a phosphoryl group from PEP to ADP is thermodynamically feasible: Notice that although the actual reaction is represented as the algebraic sum of the first two reactions, the actual reaction is a third, distinct reaction that does not involve Pi; PEP donates a phosphoryl group directly to ADP. We can describe phosphorylated compounds as having a high or low phosphoryl group transfer potential, on the basis of their standard free energies of hydrolysis (as listed in Table 13-6). The phosphoryl group transfer potential of PEP is very high, that of ATP is high, and that of glucose 6- phosphate is low (Fig. 13-19). Much of catabolism is directed toward the synthesis of high- energy phosphate compounds, but their formation is not an end in itself; they are the means of activating a wide variety of compounds for further chemical transformation. The transfer of a phosphoryl group to a compound effectively puts free energy into that compound, so that it has more free energy to give up during subsequent metabolic transformations. We described above how the synthesis of glucose 6-phosphate is accomplished by phosphoryl group transfer from ATP. In the next chapter we see how this phosphorylation of glucose activates, or “primes,” the glucose for catabolic reactions that occur in nearly every living cell. Because of its intermediate position on the scale of group transfer potential, ATP can carry energy from high-energy phosphate compounds produced by catabolism (phosphoenolpyruvate, for example) to compounds such as glucose, converting them into more reactive species with better leaving groups. ATP thus serves as the universal energy currency in all living cells. One more chemical feature of ATP is crucial to its role in metabolism: although, in aqueous solution, ATP is thermodynamically unstable and is therefore a good phosphoryl group donor, it is kinetically stable. Because of the huge activation energies (200 to 400 kJ/mol) required for uncatalyzed cleavage of its phosphoanhydride bonds, ATP does not spontaneously donate phosphoryl groups to water or to the hundreds of other potential acceptors in the cell. Only when specific enzymes are present to lower the energy of activation does phosphoryl group transfer from ATP proceed. The cell is therefore able to regulate the disposition of the energy carried by ATP by regulating the various enzymes that act on it. ATP Donates Phosphoryl, Pyrophosphoryl, and Adenylyl Groups The reactions of ATP are generally SN2 nucleophilic displacements (see Section 13.2) in which the nucleophile may be, for example, the oxygen of an alcohol or carboxylate, or a nitrogen of creatine or of the side chain of arginine or histidine. Each of the three phosphates of ATP is susceptible to nucleophilic attack (Fig. 13-20), and each position of attack yields a different type of product. FIGURE 13-20 Three positions on ATP for attack by the nucleophile R—18¨O. Any of the three P atoms (α , β , or γ ) may serve as the electrophilic target for nucleophilic attack, in this case by the labeled nucleophile R— 18Ö. The nucleophile may be an alcohol (ROH), a carboxyl group (RCOO−), or a phosphoanhydride (a nucleoside mono- or diphosphate, for example). (a) When the oxygen of the nucleophile attacks the γ position, the bridge oxygen of the product is labeled, indicating that the group transferred from ATP is a phosphoryl (—PO2−3 ), not a phosphate (—OPO2−3 ). (b) Attack on the β position displaces AMP and leads to the transfer of a pyrophosphoryl (not pyrophosphate) group to the nucleophile. (c) Attack on the α position displaces PPi and transfers the adenylyl group to the nucleophile. Nucleophilic attack by an alcohol on the γ phosphate (Fig. 13- 20a) displaces ADP and produces a new phosphate ester. Studies with 18Ö -labeled reactants have shown that the bridge oxygen in the new compound is derived from the alcohol, not from ATP; the group transferred from ATP is therefore a phosphoryl (−PO2− 3 ), not a phosphate (−OPO2− 3 ). Phosphoryl group transfer from ATP to glutamate (Fig. 13-18) or to glucose (p. 209) involves attack at the γ position of the ATP molecule. Attack at the β phosphate of ATP displaces AMP and transfers a pyrophosphoryl (not pyrophosphate) group to the attacking nucleophile (Fig. 13-20b). For example, the formation of 5- phosphoribosyl-1-pyrophosphate (Chapter 22), a key intermediate in nucleotide synthesis, results from attack of an — OH of the ribose on the β phosphate. Nucleophilic attack at the α position of ATP displaces PPi and transfers adenylate (5′-AMP) as an adenylyl group (Fig. 13-20c); the reaction is an adenylylation (a-den′-i-li-la′-shun, one of the most ungainly words in the biochemical language). Notice that hydrolysis of the α –β phosphoanhydride bond releases considerably more energy (~46 kJ/mol) than hydrolysis of the β – γ bond (~31 kJ/mol) (Table 13-6). Furthermore, the PPi formed as a byproduct of the adenylylation is hydrolyzed to two Pi by the ubiquitous enzyme inorganic pyrophosphatase, releasing 19 kJ/mol and thereby providing a further energy “push” for the adenylylation reaction. In effect, both phosphoanhydride bonds of ATP are split in the overall reaction. Adenylylation reactions are therefore thermodynamically very favorable. When the energy of ATP is used to drive a particularly unfavorable metabolic reaction, adenylylation is o en the mechanism of energy coupling. Fatty acid activation is a good example of this energy-coupling strategy. The first step in the activation of a fatty acid — either for energy- yielding oxidation or for use in the synthesis of more complex lipids — is the formation of its thiol ester (see Fig. 17-5). The direct condensation of a fatty acid with coenzyme A is endergonic, but the formation of a fatty acyl–CoA is made exergonic by stepwise removal of two phosphoryl groups from ATP. First, adenylate (AMP) is transferred from ATP to the carboxyl group of the fatty acid, forming a mixed anhydride (fatty acyl adenylate) and liberating PPi. The thiol group of coenzyme A then displaces the adenylyl group and forms a thioester with the fatty acid. The sum of these two reactions is energetically equivalent to the exergonic hydrolysis of ATP to AMP and PPi (ΔG′°=−45.6 kJ /mol) and the endergonic formation of fatty acyl–CoA. The formation of fatty acyl–CoA (ΔG′°=−31.4 kJ /mol) is made energetically favorable by hydrolysis of the PPi by inorganic pyrophosphatase. Thus, in the activation of a fatty acid, both phosphoanhydride bonds of ATP are broken. The resulting ΔG′° is the sum of the ΔG′° values for the breakage of these bonds, or −45.6 kJ /mol+ (−19.2) kJ /mol: The activation of amino acids before their polymerization into proteins (see Fig. 27-19) is accomplished by an analogous set of reactions in which a transfer RNA molecule takes the place of coenzyme A. An interesting use of the cleavage of ATP to AMP AT P + 2H2O → AM P + 2Pi ΔG′°= −64.8 kJ /mol and PPi occurs in the firefly, which uses ATP as an energy source to produce flashes of light (Box 13-2). BOX 13-2 Firefly Flashes: Glowing Reports of ATP Bioluminescence requires considerable amounts of energy. The firefly uses ATP to convert chemical energy into light energy. Males emit a flash of light to attract females, who flash in return to signal their interest. In the 1950s, from many thousands of fireflies collected by children in and around Baltimore, William McElroy and his colleagues at Johns Hopkins University isolated the principal biochemical components: luciferin, a complex carboxylic acid, and luciferase, an enzyme. Activation of luciferin by an enzymatic reaction involving pyrophosphate cleavage of ATP to form luciferyl adenylate generates the light flash (Fig. 1). In the presence of molecular oxygen and luciferase, the luciferin undergoes a multistep oxidative decarboxylation to oxyluciferin. This process is accompanied by emission of light. The color of the light flash differs from one firefly species to another and seems to be determined by differences in the structure of the luciferase. Luciferin is regenerated from oxyluciferin in a subsequent series of reactions. FIGURE 1 Important components in the firefly bioluminescence cycle. The firefly, a beetle of the Lampyridae family. In the laboratory, pure firefly luciferin and luciferase are used to measure minute quantities of ATP by the intensity of the light flash produced. As little as a few picomoles (10−12 mol) of ATP can be measured in this way. Assembly of Informational Macromolecules Requires Energy When simple precursors are assembled into high molecular weight polymers with defined sequences (DNA, RNA, proteins), as described in detail in Part III, energy is required both for the condensation of monomeric units and for the creation of ordered sequences. The precursors for DNA and RNA synthesis are nucleoside triphosphates, and polymerization is accompanied by cleavage of the phosphoanhydride linkage between the α and β phosphates, with the release of PPi (Fig. 13-20). The moieties transferred to the growing polymer in these reactions are adenylate (AMP), guanylate (GMP), cytidylate (CMP), or uridylate (UMP) for RNA synthesis, and their deoxy analogs (with TMP in place of UMP) for DNA synthesis. As noted above, the activation of amino acids for protein synthesis involves the donation of adenylyl groups from ATP, and we shall see in Chapter 27 that several steps of protein synthesis on the ribosome are also accompanied by GTP hydrolysis. In all these cases, the exergonic breakdown of a nucleoside triphosphate is coupled to the endergonic process of synthesizing a polymer of a specific sequence. ATP can supply the energy for transporting an ion or a molecule across a membrane into another aqueous compartment where its concentration is higher (see Fig. 11-39). Transport processes are major consumers of energy; in human kidney and brain, for example, as much as two-thirds of the energy consumed at rest is used to pump Na+ and K+ across plasma membranes via the Na+K+ ATPase. The transport of Na+ and K+ is driven by cyclic phosphorylation and dephosphorylation of the transporter protein, with ATP as the phosphoryl group donor. Na+-dependent phosphorylation of the Na+K+ ATPase forces a change in the protein’s conformation, and K+-dependent dephosphorylation favors return to the original conformation. Each cycle in the transport process results in the conversion of ATP to ADP and Pi, and it is the free-energy change of ATP hydrolysis that drives the cyclic changes in protein conformation that result in the electrogenic pumping of Na+ and K+. Note that in this case, ATP interacts covalently by phosphoryl group transfer to the enzyme, not to the substrate. In the contractile system of skeletal muscle cells, myosin and actin are specialized to transduce the chemical energy of ATP into motion (see Fig. 5-29). ATP binds tightly but noncovalently to one conformation of myosin, holding the protein in that conformation. When myosin catalyzes the hydrolysis of its bound ATP, the ADP and Pi dissociate from the protein, allowing it to relax into a second conformation until another molecule of ATP binds. The binding and subsequent hydrolysis of ATP (by myosin ATPase) provide the energy that forces cyclic changes in the conformation of the myosin head. The change in conformation of many individual myosin molecules results in the sliding of myosin fibrils along actin filaments (see Fig. 5-28), which translates into macroscopic contraction of the muscle fiber. As we noted earlier, this production of mechanical motion at the expense of ATP is one of the few cases in which ATP hydrolysis per se, rather than group transfer from ATP, is the source of the chemical energy in a coupled process. Transphosphorylations between Nucleotides Occur in All Cell Types Although we have focused on ATP as the cell’s energy currency and donor of phosphoryl groups, all other nucleoside triphosphates (GTP, UTP, and CTP) and all deoxynucleoside triphosphates (dATP, dGTP, dTTP, and dCTP) are energetically equivalent to ATP. The standard free-energy changes associated with hydrolysis of their phosphoanhydride linkages are very nearly identical with those shown in Table 13-6 for ATP. In preparation for their various biological roles, these other nucleotides are generated and maintained as the nucleoside triphosphate (NTP) forms by phosphoryl group transfer to the corresponding nucleoside diphosphates (NDPs) and monophosphates (NMPs). ATP is the primary high-energy phosphate compound produced by catabolism, in the processes of glycolysis, oxidative phosphorylation, and, in photosynthetic cells, photophosphorylation. Several enzymes then carry phosphoryl groups from ATP to the other nucleotides. Nucleoside diphosphate kinase, found in all cells, catalyzes the reaction 2 AT P + ND P (or dND P) M g2+ ⇌ AD P + NT P (or dNT P) ΔG′°≈ 0 Although this reaction is fully reversible, the relatively high [ATP]/[ADP] ratio in cells normally drives the reaction to the right, with the net formation of NTPs and dNTPs. The enzyme actually catalyzes a two-step phosphoryl group transfer, which is a classic case of a double-displacement (Ping-Pong) mechanism (Fig. 13- 21; see also Fig. 6-15b). First, phosphoryl group transfer from ATP to an active-site His residue produces a phosphoenzyme intermediate; then the phosphoryl group is transferred from the –His residue to an NDP acceptor. Because the enzyme is nonspecific for the base in the NDP and works equally well on dNDPs and NDPs, it can synthesize all NTPs and dNTPs, given the corresponding NDPs and a supply of ATP. FIGURE 13-21 Ping-Pong mechanism of nucleoside diphosphate kinase. The enzyme binds its first substrate (ATP in our example), and a phosphoryl group is transferred to the side chain of a His residue. ADP departs, another nucleoside (or deoxynucleoside) diphosphate replaces it, and this is converted to the corresponding triphosphate by transfer of the phosphoryl group from the phosphohistidine residue. Phosphoryl group transfers from ATP result in an accumulation of ADP; for example, when muscle is contracting vigorously, ADP accumulates and interferes with ATP-dependent contraction. During periods of intense demand for ATP, the cell lowers the ADP concentration, and at the same time replenishes ATP, by the action of adenylate kinase: 2AD P M g2+ ⇌ AT P + AM P ΔG′°≈ 0 This reaction is fully reversible, so, a er the intense demand for ATP ends, the enzyme can recycle AMP by converting it to ADP, which can then be phosphorylated to ATP in mitochondria. A similar enzyme, guanylate kinase, converts GMP to GDP at the expense of ATP. By pathways such as these, energy conserved in the catabolic production of ATP is used to supply the cell with all required NTPs and dNTPs. Phosphocreatine (PCr; Fig. 13-15), also called creatine phosphate, serves as a ready source of phosphoryl groups for the quick synthesis of ATP from ADP. The PCr concentration in skeletal muscle is approximately 30 mM, nearly 10 times the concentration of ATP, and in other tissues such as smooth muscle, brain, and kidney, [PCr] is 5 to 10 mM. The enzyme creatine kinase catalyzes the reversible reaction AD P + PCr M g2+ ⇌ AT P + Cr ΔG′°= −12.5 kJ /mol When a sudden demand for energy depletes ATP, the PCr reservoir is used to replenish ATP at a rate considerably faster than ATP can be synthesized by catabolic pathways. When the demand for energy slackens, ATP produced by catabolism is used to replenish the PCr reservoir by reversal of the creatine kinase reaction (see Box 23-1). Organisms in the lower phyla employ other PCr-like molecules (collectively called phosphagens) as phosphoryl reservoirs. SUMMARY 13.3 Phosphoryl Group Transfers and ATP ATP is the chemical link between catabolism and anabolism. It is the energy currency of the living cell. The exergonic conversion of ATP to ADP and Pi, or to AMP and PPi, is coupled to many endergonic reactions and processes. The free-energy change for ATP hydrolysis under cellular conditions is its phosphorylation potential, ΔGp. Direct hydrolysis of ATP is the source of energy in some processes driven by conformational changes. In general, however, it is not ATP hydrolysis but the transfer of a phosphoryl group from ATP to a substrate or an enzyme that couples the energy of ATP breakdown to endergonic transformations of substrates. Phosphate compounds with high free energies of hydrolysis can donate their phosphoryl group to form another phosphate compound with a smaller free energy of hydrolysis. ATP can also donate a pyrophosphoryl (PPi) or adenylyl (AMP) group to a variety of metabolic intermediates, activating them for nucleophilic displacement reactions. Through these group transfer reactions, ATP provides the energy for a large number of anabolic reactions, including the synthesis of informational macromolecules, and for the transport of molecules and ions across membranes against concentration gradients and electrical potential gradients. Muscle contraction is also powered by ATP. To maintain its high group transfer potential, ATP concentration must be held far above the equilibrium concentration by energy-yielding reactions of catabolism. ATP can donate a phosphoryl group to nucleoside diphosphates by transphosphorylation to keep the levels of GTP, UTP, CTP, and the deoxynucleotides far above their equilibrium concentrations. 13.4 Biological Oxidation- Reduction Reactions The transfer of phosphoryl groups is a central feature of metabolism. Equally important is another kind of transfer: electron transfer in oxidation-reduction reactions, sometimes referred to as redox reactions. These reactions involve the loss of electrons by one chemical species, which is thereby oxidized, and the gain of electrons by another, which is reduced. The flow of electrons in oxidation-reduction reactions is responsible, directly or indirectly, for all work done by living organisms. In nonphotosynthetic organisms, the sources of electrons are reduced compounds (foods); in photosynthetic organisms, the initial electron donor is a chemical species excited by the absorption of light. The path of electron flow in metabolism is complex. Electrons move from various metabolic intermediates to specialized electron carriers in enzyme-catalyzed reactions. The carriers, in turn, donate electrons to acceptors with higher electron affinities, with the release of energy. Cells possess a variety of molecular energy transducers, which convert the energy of electron flow into useful work. We begin by discussing how work can be accomplished by an electromotive force (emf), then consider the theoretical and experimental basis for measuring energy changes in oxidation reactions in terms of emf and the relationship between this force, expressed in volts, and the free-energy change, expressed in joules. We also describe the structures and oxidation-reduction chemistry of the most common of the specialized electron carriers, which you will encounter repeatedly in later chapters. The Flow of Electrons Can Do Biological Work Every time we use a motor, an electric light or heater, or a spark to ignite gasoline in a car engine, we use the flow of electrons to accomplish work. In the circuit that powers a motor, the source of electrons can be a battery containing two chemical species that differ in affinity for electrons. Electrical wires provide a pathway for electron flow from the chemical species at one pole of the battery, through the motor, to the chemical species at the other pole of the battery. Because the two chemical species differ in their affinity for electrons, electrons flow spontaneously through the circuit, driven by a force proportional to the difference in electron affinity, the electromotive force, emf. The emf (typically a few volts) can accomplish work if an appropriate energy transducer — in this case a motor — is placed in the circuit. The motor can be coupled to a variety of mechanical devices to do useful work. Living cells have an analogous biological “circuit,” with a relatively reduced compound such as glucose as the source of electrons. As glucose is enzymatically oxidized, the released electrons flow spontaneously through a series of electron-carrier intermediates to another chemical species, such as O2. This electron flow is exergonic, because O2 has a higher affinity for electrons than do the electron-carrier intermediates. The resulting emf provides energy to a variety of molecular energy transducers (enzymes and other proteins) that do biological work. In the mitochondrion, for example, membrane-bound enzymes couple electron flow to the production of a transmembrane pH difference and a transmembrane electrical potential, accomplishing chemiosmotic and electrical work. The proton gradient thus formed has potential energy, sometimes called the proton-motive force by analogy with electromotive force. Another enzyme, ATP synthase in the inner mitochondrial membrane, uses the proton-motive force to do chemical work: synthesis of ATP from ADP and Pi as protons flow spontaneously across the membrane. Similarly, membrane-localized enzymes in Escherichia coli convert emf to proton-motive force, which is then used to power flagellar motion. The principles of electrochemistry that govern energy changes in the macroscopic circuit with a motor and battery apply with equal validity to the molecular processes accompanying electron flow in living cells. Oxidation-Reductions Can Be Described as Half-Reactions Although oxidation and reduction must occur together, it is convenient when describing electron transfers to consider the two halves of an oxidation-reduction reaction separately. For example, the oxidation of ferrous ion by cupric ion, Fe2+ + Cu2+ ⇌ Fe3+ + Cu+ can be described in terms of two half-reactions: 1. Fe2+ ⇌ Fe3+ + e− 2. Cu2+ + e− ⇌ Cu+ The electron-donating molecule in an oxidation-reduction reaction is called the reducing agent or reductant; the electron- accepting molecule is the oxidizing agent or oxidant. A given agent, such as an iron cation existing in the ferrous (Fe2+) or ferric (Fe3+) state, functions as a conjugate reductant-oxidant pair (redox pair), just as an acid and corresponding base function as a conjugate acid-base pair. Recall from Chapter 2 that in acid- base reactions we can write a general equation: proton donor ⇌ H++ proton acceptor. In redox reactions we can write a similar general equation: electron donor (reductant) ⇌ e−+ electron acceptor (oxidant). In the reversible half-reaction (1) above, Fe2+ is the electron donor and Fe3+ is the electron acceptor; together, Fe2+ and Fe3+ constitute a conjugate redox pair. The mnemonic OIL RIG — oxidation is losing, reduction is gaining — may be helpful in remembering what happens to electrons in redox reactions. The electron transfers in the oxidation-reduction reactions of organic compounds are not fundamentally different from those of inorganic species. Consider the oxidation of a reducing sugar (an aldehyde or a ketone) by cupric ion: This overall reaction can be expressed as two half-reactions: 1. 2. 2Cu2+ + 2e− + 2OH− ⇌ Cu2O + H2O Notice that because two electrons are removed from the aldehyde carbon, the second half-reaction (the one-electron reduction of cupric to cuprous ion) must be doubled to balance the overall equation. Biological Oxidations O en Involve Dehydrogenation The carbon in living cells exists in a range of oxidation states (Fig. 13-22). When a carbon atom shares an electron pair with another atom (typically H, C, S, N, or O), the sharing is unequal, in favor of the more electronegative atom. The order of increasing electronegativity is H < C < S < N< O. In oversimplified but useful terms, the more electronegative atom “owns” the bonding electrons it shares with another atom. For example, in methane (CH4), carbon is more electronegative than the four hydrogens bonded to it, and the C atom therefore owns all eight bonding electrons (Fig. 13-22). In ethane, the electrons in the C— C bond are shared equally, so each C atom owns only seven of its eight bonding electrons. In ethanol, C-1 is less electronegative than the oxygen to which it is bonded, and the O atom therefore owns both electrons of the C— O bond, leaving C-1 with only five bonding electrons. With each formal loss of “owned” electrons, the carbon atom has undergone oxidation — even when no oxygen is involved, as in the conversion of an alkane (—CH2—CH2—) to an alkene (—CH═CH—). In this case, oxidation (loss of electrons) is coincident with the loss of hydrogen. In biological systems, as we noted earlier in the chapter, oxidation is o en synonymous with dehydrogenation, and many enzymes that catalyze oxidation reactions are dehydrogenases. Notice that the more reduced compounds in Figure 13-22 (top) are richer in hydrogen than in oxygen, whereas the more oxidized compounds (bottom) have more oxygen and less hydrogen.
FIGURE 13-22 Different levels of oxidation of carbon compounds in the biosphere. To approximate the level of oxidation of these compounds, focus on the red carbon atom and its bonding electrons. When this carbon is bonded to the less electronegative H atom, both bonding electrons (red) are assigned to the carbon. When carbon is bonded to another carbon, bonding electrons are shared equally, so one of the two electrons is assigned to the red carbon. When the red carbon is bonded to the more electronegative O atom, the bonding electrons are assigned to the oxygen. The number to the right of each compound is the number of electrons “owned” by the red carbon, a rough expression of the degree of oxidation of that compound. As the red carbon undergoes oxidation (loses electrons), the number gets smaller. Not all biological oxidation-reduction reactions involve carbon. For example, in the conversion of molecular nitrogen to ammonia, 6H+ + 6e− + N2 → 2NH3, the nitrogen atoms are reduced. Electrons are transferred from one molecule (electron donor) to another (electron acceptor) in one of four ways: 1. Directly as electrons. For example, the Fe2+/Fe3+ redox pair can transfer an electron to the Cu+/Cu2+ redox pair: Fe2+ + Cu2+ ⇌ Fe3+ + Cu+ 2. As hydrogen atoms. Recall that a hydrogen atom consists of a proton (H+) and a single electron (e−). In this case we can write the general equation AH2 ⇌ A + 2e− + 2H+ where AH2 is the hydrogen/electron donor. (Do not mistake the above reaction for an acid dissociation, which involves a proton and no electron.) AH2 and A together constitute a conjugate redox pair (A/AH2), which can reduce another compound B (or redox pair, B/BH2) by transfer of hydrogen atoms: AH2+ B ⇌ A + BH2 3. As a hydride ion (:H−), which has two electrons. This occurs in the case of NAD-linked dehydrogenases, described below. 4. Through direct combination with oxygen. In this case, oxygen combines with an organic reductant and is covalently incorporated in the product, as in the oxidation of a hydrocarbon to an alcohol: R—CH3+ O2 → R—CH2—OH The hydrocarbon is the electron donor, and the oxygen atom is the electron acceptor. All four types of electron transfer occur in cells. The neutral term reducing equivalent is commonly used to designate a single electron equivalent participating in an oxidation-reduction reaction, no matter whether this equivalent is an electron per se or is part of a hydrogen atom or a hydride ion, or whether the electron transfer takes place in a reaction with oxygen to yield an oxygenated product. 1 2 Reduction Potentials Measure Affinity for Electrons When two conjugate redox pairs are together in solution, electron transfer from the electron donor of one pair to the electron acceptor of the other may proceed spontaneously. The tendency for such a reaction depends on the relative affinity of the electron acceptor of each redox pair for electrons. The standard reduction potential, E°, a measure (in volts) of this affinity, can be determined in an experiment such as that described in Figure 13- 23. Electrochemists have chosen as a standard of reference the half-reaction H+ + e− → H2 The electrode at which this half-reaction occurs (called a half- cell) is arbitrarily assigned an E° of 0.00 V. When this hydrogen electrode is connected through an external circuit to another half- cell in which an oxidized species and its corresponding reduced species are present at standard concentrations (at 25 °C, each solute at 1 M, each gas at 101.3 kPa), electrons tend to flow through the external circuit from the half-cell of lower E° to the half-cell of higher E°. By convention, a half-cell that takes electrons from the standard hydrogen cell is assigned a positive value of E°, and one that donates electrons to the hydrogen cell, a negative value. When any two half-cells are connected, that with 1 2 the larger (more positive) E° will be reduced; it has the greater reduction potential. FIGURE 13-23 Measurement of the standard reduction potential (E′°) of a redox pair. Electrons flow from the test electrode to the reference electrode, or vice versa. The ultimate reference half-cell is the hydrogen electrode, as shown here, at pH 0. The electromotive force (emf) of this electrode is designated 0.00 V. At pH 7 in the test cell (at 25 °C), E′° for the hydrogen electrode is −0.414 V. The direction of electron flow depends on the relative electron “pressure” or potential of the two cells. A salt bridge containing a saturated KCl solution provides a path for counter-ion movement between the test cell and the reference cell. From the observed emf and the known emf of the reference cell, the experimenter can find the emf of the test cell containing the redox pair. The cell that gains electrons has, by convention, the more positive reduction potential. The reduction potential of a half-cell depends not only on the chemical species present but also on their activities, approximated by their concentrations. The Nernst equation relates standard reduction potential (E°) to the actual reduction potential (E) at any concentration of oxidized and reduced species in a living cell: E = E°+ ln (13-5) where R and T have their usual meanings, n is the number of electrons transferred per molecule, and F is the Faraday constant, a proportionality constant that converts volts to joules (Table 13- 1). At 298 K (25 °C), this expression reduces to E = E°+ ln (13-6) RT nF [electron acceptor] [electron donor] 0.026 V n [electron acceptor] [electron donor] KEY CONVENTION Many half-reactions of interest to biochemists involve protons. As in the definition of ΔG ′°, biochemists define the standard state for oxidation-reduction reactions as pH 7 and express a standard transformed reduction potential, E′°, the standard reduction potential at pH 7 and 25 °C. By convention, ΔE′° for any redox reaction is given as E′° of the electron acceptor minus E′° of the electron donor. The standard reduction potentials given in Table 13-7 and used throughout this book are values for E′° and are therefore valid only for systems at neutral pH. Each value represents the potential difference when the conjugate redox pair, at 1 M concentrations, 25 °C, and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2H+/H2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the standard (pH 0) cell; the measured E′° for the 2H+/H2 pair is –0.414 V. TABLE 13-7 Standard Reduction Potentials of Some Biologically Important Half-Reactions Half-reaction E′° (V) O2+ 2H+ + 2e− → H2O 0.8161 2 Fe3+ + e− → Fe2+ 0.771 NO− 3+ 2H+ + 2e− → NO− 2 + H2O 0.421 Cytochrome f (Fe3+)+ e− → cytochrome f (Fe2+) 0.365 Fe(CN)3− 6 (ferricyanide)+ e− → Fe(CN)4− 6 0.36 Cytochrome a3 (Fe3+)+ e− → cytochrome a3 (Fe2+) 0.35 O2+ 2H+ + 2e− → H2O2 0.295 Cytochrome a (Fe3+)+ e− → cytochrome a (Fe2+) 0.29 Cytochrome c (Fe3+)+ e− → cytochrome c (Fe2+) 0.254 Cytochrome c1 (Fe3+)+ e− → cytochrome c1 (Fe2+) 0.22 Cytochrome b (Fe3+)+ e− → cytochrome b (Fe2+) 0.077 U biquinone+ 2H+ + 2e− → ubiquinol 0.045 Fumarate2− + 2H+ + 2e− → succinate2− 0.031 2H+ + 2e− → H2 (at standard conditions, pH 0) 0.000 Crotonyl-CoA + 2H+ + 2e− → butyryl-CoA −0.015 Oxaloacetate2− + 2H+ + 2e− → malate2− −0.166 Pyruvate− + 2H+ + 2e− → lactate− −0.185 Acetaldehyde+ 2H+ + 2e− → ethanol −0.197 FAD + 2H+ + 2e− → FAD H2 −0.219 G lutathione+ 2H+ + 2e− → 2 reduced glutathione −0.23 S + 2H+ + 2e− → H2S −0.243 a Lipoic acid+ 2H+ + 2e− → dihydrolipoic acid −0.29 NAD + + H+ + 2e− → NAD H −0.320 NAD P+ + H+ + 2e− → NAD PH −0.324 Acetoacetate+ 2H+ + 2e− → β-hydroxybutyrate −0.346 α-Ketoglutarate+ CO2+ 2H+ + 2e− → isocitrate −0.38 2H+ + 2e− → H2 (at pH 7) −0.414 Ferredoxin (Fe3+)+ e− → ferredoxin (Fe2+) −0.432 Data mostly from R. A. Loach, in Handbook of Biochemistry and Molecular Biology, 3rd edn (G. D. Fasman, ed.), Physical and Chemical Data, Vol. 1, p. 122, CRC Press, 1976. This is the value for free FAD; FAD bound to a specific flavoprotein (e.g., succinate dehydrogenase) has a different E′° that depends on its protein environment. Standard Reduction Potentials Can Be Used to Calculate Free-Energy Change Why are reduction potentials so useful to the biochemist? When E values have been determined for any two half-cells, relative to the standard hydrogen electrode, we also know their reduction potentials relative to each other. We can then predict the direction in which electrons will tend to flow when the two half-cells are connected through an external circuit or when components of both half-cells are present in the same solution. Electrons tend to flow to the half-cell with the more positive E, and the strength of a that tendency is proportional to ∆E, the difference in reduction potential. The energy made available by this spontaneous electron flow (the free-energy change, ∆G, for the oxidation- reduction reaction) is proportional to ∆E: ΔG = −nFΔE or ΔG ′°= −nFΔE′° (13-7) where n is the number of electrons transferred in the reaction. With this equation we can calculate the actual free-energy change for any oxidation-reduction reaction from the values of ΔE′° in a table of reduction potentials (Table 13-7) and the concentrations of reacting species. WORKED EXAMPLE 13-3 Calculation of ΔG ′° and Δ G of a Redox Reaction Calculate the standard free-energy change, ΔG ′°, for the reaction in which acetaldehyde is reduced by the biological electron carrier NADH: Acetaldehyde+ NAD H + H+ → ethanol+ NAD + Then calculate the actual free-energy change, ∆G, when [acetaldehyde] and [NADH] are 1.00 M, and [ethanol] and [NAD +] are 0.100 M. The relevant half-reactions and their E′° values are 1. Acetaldehyde+ 2H+ + 2e− → ethanol E′°= −0.197 V 2. NAD + + 2H+ + 2e− → NAD H + H+ E′°= −0.320 V Remember that, by convention, ΔE′° is E′° of the electron acceptor minus E′° of the electron donor. It represents the difference between the electron affinities of the two half- reactions in the table of reduction potentials (Table 13-7). Note that the more widely separated the two half-reactions in the table, the more energetic the electron-transfer reaction when the two half-reactions occur together. By convention, in tables of reduction potentials, all half-reactions are represented as reductions, but when two half-reactions occur together, one of them must be an oxidation. Although that half-reaction will go in the opposite direction from that shown in Table 13-7, we do not change the sign of that half-reaction before calculating ΔE′°, because ΔE′° is defined as a difference of reduction potentials. SOLUTION: Because acetaldehyde is accepting electrons (n= 2) from NADH, ΔE′°=−0.197 V − (−0.320 V)= 0.123 V. Therefore, ΔG ′°=−nFΔE′°=−2(96.5 kJ /V ∙mol) (0.123 V)=−23.7 kJ /mo This is the free-energy change for the oxidation-reduction reaction at 25 °C and pH 7, when acetaldehyde, ethanol, NAD +, and NADH are all present at 1.00 M concentrations. To calculate Δ G when [acetaldehyde] and [NADH] are 1.00 M, and [ethanol] and [NAD +] are 0.100 M, we can use Equation 13-4 and the standard free-energy change calculated above: This is the actual free-energy change at the specified concentrations of the redox pairs. A Few Types of Coenzymes and Proteins Serve as Universal Electron Carriers The principles of oxidation-reduction energetics described above apply to the many metabolic reactions that involve electron ΔG = ΔG ′°+ RT ln = −23.7 kJ /mol+ (8.315 J /mol∙K)(298 K) ln = −23.7 kJ /mol+ (2.48 J /mol) ln 0.01 = −35.1 kJ /mol [ethanol][NAD +] [acetaldehyde][NAD H] (0.100 M )( (1.00 M )( transfers. For example, in many organisms, the oxidation of glucose supplies energy for the production of ATP. The complete oxidation of glucose C6H12O6+ 6O2 → 6CO2+ 6H2O has a ΔG ′° of −2,840 kJ /mol. This is a much larger release of free energy than is required for ATP synthesis in cells (50 to 60 kJ/mol; see Worked Example 13-2). Cells convert glucose to CO2 not in a single, high-energy-releasing reaction but rather in a series of controlled reactions, some of which are oxidations. The free energy released in these oxidation steps is of the same order of magnitude as that required for ATP synthesis from ADP, with some energy to spare. Electrons removed in these oxidation steps are transferred to coenzymes specialized for carrying electrons, such as NAD + and FAD (described below). The multitude of enzymes that catalyze cellular oxidations channel electrons from their hundreds of different substrates into just a few types of universal electron carriers. The reduction of these carriers in catabolic processes results in the conservation of free energy released by substrate oxidation. NAD, NADP, FMN, and FAD are water-soluble coenzymes that undergo reversible oxidation and reduction in many of the electron-transfer reactions of metabolism. The nucleotides NAD and NADP move readily from one enzyme to another; the flavin nucleotides FMN and FAD are usually very tightly bound to the enzymes, called flavoproteins, for which they serve as prosthetic groups. Lipid- soluble quinones such as ubiquinone and plastoquinone act as electron carriers and proton donors in the nonaqueous environment of membranes. Iron-sulfur proteins and cytochromes, which have tightly bound prosthetic groups that undergo reversible oxidation and reduction, also serve as electron carriers in many oxidation-reduction reactions. Some of these proteins are water-soluble, but others are peripheral or integral membrane proteins. The oxidation-reduction chemistry of quinones, iron-sulfur proteins, and cytochromes is discussed in Chapters 19 and 20. Nicotinamide adenine dinucleotide (NAD; NAD + in its oxidized form) and its close analog nicotinamide adenine dinucleotide phosphate (NADP; NAD P+ when oxidized) are composed of two nucleotides joined through their phosphate groups by a phosphoanhydride bond (Fig. 13-24a). Because the nicotinamide ring resembles pyridine, these compounds are sometimes called pyridine nucleotides. The vitamin niacin is the source of the nicotinamide moiety in nicotinamide nucleotides. FIGURE 13-24 NAD and NADP. (a) Nicotinamide adenine dinucleotide, NAD +, and its phosphorylated analog, NAD P+, undergo reduction to NADH and NADPH, accepting a hydride ion (two electrons and one proton) from an oxidizable substrate. The hydride ion is added to either the front or the back of the planar nicotinamide ring. (b) The UV absorption spectra of NAD + and NADH. Reduction of the nicotinamide ring produces a new, broad absorption band with a maximum at 340 nm. The production of NADH during an enzyme-catalyzed reaction can be conveniently followed by observing the appearance of the absorbance at 340 nm (molar extinction coefficient ε340= 6,200 M −1 cm−1). Both coenzymes undergo reversible reduction of the nicotinamide ring (Fig. 13-24). As a substrate molecule undergoes oxidation (dehydrogenation), giving up two hydrogen atoms, the oxidized form of the nucleotide (NAD + or NAD P+) accepts a hydride ion (:H−, the equivalent of a proton and two electrons) and is reduced (to NADH or NADPH). The second proton removed from the substrate is released to the aqueous solvent. The half- reactions for these nucleotide cofactors are NAD + + 2e− + 2H+ → NAD H + H+ NAD P+ + 2e− + 2H+ → NAD PH + H+ Reduction of NAD + or NAD P+ converts the benzenoid ring of the nicotinamide moiety (with a fixed positive charge on the ring nitrogen) to the quinonoid form (with no charge on the nitrogen). The reduced nucleotides absorb light at 340 nm; the oxidized forms do not (Fig. 13-24b). Biochemists use this difference in absorption to assay reactions involving these coenzymes. Note that the plus sign in the abbreviations NAD + and NAD P+ does not indicate the net charge on these molecules (in fact, both are negatively charged); rather, it indicates that the nicotinamide ring is in its oxidized form, with a positive charge on the nitrogen atom. In the abbreviations NADH and NADPH, the “H” denotes the added hydride ion. To refer to these nucleotides without specifying their oxidation state, we use NAD and NADP. The total concentration of NAD + + NAD H in most tissues is about 10−5M ; that of NAD P+ + NAD PH is about 10−6M . In many cells and tissues, the ratio of NAD + (oxidized) to NADH (reduced) is high, favoring hydride transfer from a substrate to NAD + to form NADH. By contrast, NADPH is generally present at a higher concentration than NAD P+, favoring hydride transfer from NADPH to a substrate. This reflects the specialized metabolic roles of the two coenzymes: NAD + generally functions in oxidations — usually as part of a catabolic reaction; NADPH is the usual coenzyme in reductions — nearly always as part of an anabolic reaction. A few enzymes can use either coenzyme, but most show a strong preference for one over the other. Also, the processes in which these two cofactors function are segregated in eukaryotic cells: for example, oxidations of fuels such as pyruvate, fatty acids, and α -keto acids derived from amino acids occur in the mitochondrial matrix, whereas reductive biosynthetic processes such as fatty acid synthesis take place in the cytosol. This functional and spatial specialization allows a cell to maintain two distinct pools of electron carriers, with two distinct functions. More than 200 enzymes are known to catalyze reactions in which NAD + (or NAD P+) accepts a hydride ion from a reduced substrate, or NADPH (or NADH) donates a hydride ion to an oxidized substrate. The general reactions are AH2+ NAD + → A + NAD H + H+ A + NAD PH + H+ → AH2+ NAD P+ where AH2 is the reduced substrate and A is the oxidized substrate. The general name for an enzyme of this type is oxidoreductase; they are also commonly called dehydrogenases. For example, alcohol dehydrogenase catalyzes the first step in the catabolism of ethanol, in which ethanol is oxidized to acetaldehyde: CH3CH2OH Ethanol + NAD + → CH3CHO + NAD H + H+ Acetaldehyde Notice that one of the carbon atoms in ethanol has lost a hydrogen; the compound has been oxidized from an alcohol to an aldehyde (refer again to Fig. 13-22 for the oxidation states of carbon). The association between a dehydrogenase and NAD or NADP is relatively loose; the coenzyme readily diffuses from one enzyme to another, acting as a water-soluble carrier of electrons from one metabolite to another. For example, in the production of alcohol during fermentation of glucose by yeast cells, a hydride ion is removed from glyceraldehyde 3-phosphate by one enzyme (glyceraldehyde 3-phosphate dehydrogenase) and transferred to NAD +. The NADH produced then leaves the enzyme surface and diffuses to another enzyme (alcohol dehydrogenase), which transfers a hydride ion to acetaldehyde, producing ethanol: (1) G lyceraldehyde 3-phosphate+ NAD + → 3-phosphoglycerate+ NAD H + H+ (2) Acetaldehyde+ NAD H + H+ → ethanol+ NAD + –––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– Sum :G lyceraldehyde 3-phosphate+ acetaldehyde→ 3-phosphoglycerate+ ethanol Notice that in the overall reaction there is no net production or consumption of NAD + or NADH; the coenzymes function catalytically and are recycled repeatedly without a net change in the total amount of NAD + + NAD H. Both reduced and oxidized forms of NAD and NADP serve as allosteric effectors of proteins in catabolic pathways. As we describe in later chapters, the ratios NAD +/NAD H and NAD P+/NAD PH serve as sensitive gauges of a cell’s fuel supply, allowing rapid, appropriate changes in energy-yielding and energy-dependent metabolism. NAD Has Important Functions in Addition to Electron Transfer Some key cellular functions are regulated by enzymes that use NAD + not as a redox cofactor but as a substrate in a coupled reaction in which the availability of NAD + can be an indicator of the cell’s energy status. In DNA replication and repair, the enzyme DNA ligase is adenylylated and then transfers the AMP to a 5′ phosphate in a nicked DNA (see Fig. 25-15); in bacteria, NAD + serves as the source of the activating AMP group. A family of proteins called sirtuins regulate the activity of proteins in diverse cellular pathways by deacetylating the ε-amino group of an acetylated Lys residue. The deacetylation is coupled to NAD + hydrolysis, yielding O-acetyl-ADP-ribose and nicotinamide. Among the cellular processes regulated by sirtuins are inflammation, apoptosis, aging, and DNA transcription; deacetylation by a sirtuin alters the charge on histones, influencing which genes are expressed. The availability of NAD + for these types of reactions may indicate that the cell is undergoing stress and that pathways designed to respond to stress should be activated. NAD + also plays an important role in cholera pathology (see Section 12.2). Cholera toxin has an enzymatic activity that transfers ADP-ribose from NAD + to a G protein involved in regulating ion fluxes in the cells lining the gut. This ADP- ribosylation blocks water retention, causing the diarrhea and dehydration characteristic of cholera. Dietary deficiency of niacin, the vitamin form of NAD and NADP, causes pellagra (Fig. 13-25). The pyridine-like rings of NAD and NADP are derived from the vitamin niacin (nicotinic acid; Fig. 13- 26), which is synthesized from tryptophan. Humans generally cannot synthesize sufficient quantities of niacin, and this is especially so for individuals with diets low in tryptophan (maize, for example, has a low tryptophan content). Niacin deficiency, which affects all the NAD(P)-dependent dehydrogenases, causes the serious human disease pellagra (Italian for “rough skin”) and a related disease in dogs, called black tongue. Pellagra is characterized by the “three Ds”: dermatitis, diarrhea, and dementia, followed in many cases by death. A century ago, pellagra was a common human disease; in the southern United States, where maize was a dietary staple, about 100,000 people were afflicted and about 10,000 died as a result of this disease between 1912 and 1916. In 1920, Joseph Goldberger showed pellagra to be caused by a dietary insufficiency, and in 1937, Frank Strong, D. Wayne Woolley, and Conrad Elvehjem identified niacin as the curative agent for the dog version of pellagra, black tongue. Supplementation of the human diet with this inexpensive compound has nearly eradicated pellagra in the populations of the developed world, with one significant exception: people who drink excessive amounts of alcohol. In these individuals, intestinal absorption of niacin is much reduced, and caloric needs are o en met with distilled spirits that are virtually devoid of vitamins, including niacin.
FIGURE 13-25 Dermatitis associated with pellagra. Dermatitis involving the face, hands, and feet is an early sign of pellagra, a serious human disease that results from insufficient niacin in the diet. Untreated, pellagra leads to dementia and ultimately is fatal. FIGURE 13-26 Niacin (nicotinic acid) and its derivative nicotinamide. The biosynthetic precursor of these compounds is tryptophan. In the laboratory, nicotinic acid was first produced by oxidation of the natural product nicotine—thus the name. Both nicotinic acid and nicotinamide cure pellagra, but nicotine (from cigarettes or elsewhere) has no curative activity. Flavin Nucleotides Are Tightly Bound in Flavoproteins Flavoproteins are enzymes that catalyze oxidation-reduction reactions using either flavin mononucleotide (FMN) or flavin adenine dinucleotide (FAD) as coenzyme (Fig. 13-27). These coenzymes, the flavin nucleotides, are derived from the vitamin riboflavin. The fused ring structure of flavin nucleotides (the isoalloxazine ring) undergoes reversible reduction, accepting either one or two electrons in the form of one or two hydrogen atoms (each atom an electron plus a proton) from a reduced substrate. The fully reduced forms are abbreviated FAD H2 and FM NH2. When a fully oxidized flavin nucleotide accepts only one electron (one hydrogen atom), the semiquinone form of the isoalloxazine ring is produced, abbreviated FAD H∙ and FM NH∙. Because flavin nucleotides have a chemical specialty that is slightly different from that of the nicotinamide coenzymes — the ability to participate in either one- or two-electron transfers — flavoproteins are involved in a greater diversity of reactions than the NAD(P)-linked dehydrogenases. FIGURE 13-27 Oxidized and reduced FAD and FMN. FMN consists of the structure above the dashed red line across the FAD molecule (oxidized form). The flavin nucleotides accept two hydrogen atoms (two electrons and two protons), both of which appear in the flavin ring system (isoalloxazine ring). When FAD or FMN accepts only one hydrogen atom, the semiquinone, a stable free radical, forms. Like the nicotinamide coenzymes (Fig. 13-24), the flavin nucleotides undergo a shi in a major absorption band on reduction (again, useful to biochemists who want to monitor reactions involving these coenzymes). Flavoproteins that are fully reduced (two electrons accepted) generally have an absorption maximum near 360 nm. When partially reduced (one electron), they acquire another absorption maximum at about 450 nm; when fully oxidized, the flavin has maxima at 370 and 440 nm. The flavin nucleotide in most flavoproteins is bound rather tightly to the protein, and in some enzymes, such as succinate dehydrogenase, it is bound covalently. Such tightly bound coenzymes are properly called prosthetic groups. They do not transfer electrons by diffusing from one enzyme to another; rather, they provide a means by which the flavoprotein can temporarily hold electrons while it catalyzes electron transfer from a reduced substrate to an electron acceptor. One important feature of the flavoproteins is the variability in the standard reduction potential (E′°) of the bound flavin nucleotide. Tight association between the enzyme and prosthetic group confers on the flavin ring a reduction potential typical of that particular flavoprotein, sometimes quite different from the reduction potential of the free flavin nucleotide. FAD bound to succinate dehydrogenase, for example, has an E′° close to 0.0 V, compared with −0.219 V for free FAD; E′° for other flavoproteins ranges from −0.40 V to+ 0.06 V. Flavoproteins are o en very complex; some have, in addition to a flavin nucleotide, tightly bound inorganic ions (iron or molybdenum, for example) capable of participating in electron transfers. We examine the function of flavoproteins as electron carriers in Chapters 19 and 20, when we consider their roles in oxidative phosphorylation (in mitochondria) and photophosphorylation (in chloroplasts). SUMMARY 13.4 Biological Oxidation- Reduction Reactions In many organisms, a central energy-conserving process is the stepwise oxidation of glucose to CO2, in which some of the energy of oxidation is conserved in ATP as electrons are passed to O2. Biological oxidation-reduction reactions can be described in terms of two half-reactions, each with a characteristic standard reduction potential, E′°. Many biological oxidation reactions are dehydrogenations in which one or two hydrogen atoms (H+ + e−) are transferred from a substrate to a hydrogen acceptor. In some biological redox reactions, the substrate loses both electrons and protons, the equivalent of losing hydrogen. The many enzymes that catalyze such reactions are called dehydrogenases. When two electrochemical half-cells, each containing the components of a half-reaction, are connected, electrons tend to flow to the half-cell with the higher reduction potential. The strength of this tendency is proportional to the difference between the two reduction potentials (∆E) and is a function of the concentrations of oxidized and reduced species. The standard free-energy change for an oxidation-reduction reaction is directly proportional to the difference in standard reduction potentials of the two half-cells: ΔG ′°= −nFΔE′°. Oxidation-reduction reactions in living cells involve specialized electron carriers. NAD and NADP are the freely diffusible coenzymes of many dehydrogenases. Both NAD + and NAD P+ accept two electrons and one proton. In addition to its role in oxidation-reduction reactions, NAD + is the source of AMP in the bacterial DNA ligase reaction and of ADP-ribose in the cholera toxin reaction, and it is hydrolyzed in the deacetylation of proteins by some sirtuins. Lack of the vitamin niacin prevents NAD synthesis and leads to pellagra. FAD and FMN, the flavin nucleotides, serve as tightly bound prosthetic groups of flavoproteins. They can accept either one or two electrons and one or two protons. Their reduction potentials depend on the flavoprotein with which they are associated. 13.5 Regulation of Metabolic Pathways Metabolic regulation is one of the most remarkable features of living organisms. Of the thousands of enzyme-catalyzed reactions that can take place in a cell, there is probably not one that escapes some form of regulation. This need to regulate every aspect of cellular metabolism becomes clear as one examines the complexity of metabolic reaction sequences. Although it is convenient for the student of biochemistry to divide metabolic processes into “pathways” that play discrete roles in the cell’s economy, no such separation exists in the living cell. Rather, every pathway we discuss in this book is inextricably intertwined with all the other cellular pathways in a multidimensional network of reactions (Fig. 13-28).
FIGURE 13-28 Metabolism as a three-dimensional meshwork. A typical eukaryotic cell has the capacity to make about 30,000 different proteins, which catalyze thousands of different reactions involving many hundreds of metabolites, most shared by more than one “pathway.” In this much- simplified overview of metabolic pathways, each dot represents an intermediate compound and each connecting line represents an enzymatic reaction. For a more realistic and far more complex diagram of metabolism, see the online KEGG PATHWAY database (www.genome.ad.jp/kegg/pathway/map/map01100.html); in this interactive map, you can click on each dot to obtain extensive data about the compound and the enzymes for which it is a substrate. For example, in Chapter 14 we will discuss four possible fates for glucose 6- phosphate in a hepatocyte: breakdown by glycolysis for the production of ATP, breakdown in the pentose phosphate pathway for the production of NADPH and pentose phosphates, use in the synthesis of complex polysaccharides of the extracellular matrix, or hydrolysis to glucose and phosphate to replenish blood glucose. In fact, glucose 6-phosphate has other possible fates in hepatocytes, too; it may, for example, be used to synthesize other sugars, such as glucosamine, galactose, galactosamine, fucose, and neuraminic acid, for use in protein glycosylation, or it may be partially degraded to provide acetyl-CoA for fatty acid and sterol synthesis. And E. coli can use glucose to produce the carbon skeleton of every one of its several thousand types of molecules. When any cell uses glucose 6- phosphate for one purpose, that “decision” affects all the other pathways for which glucose 6-phosphate is a precursor or intermediate: any change in the allocation of glucose 6-phosphate to one pathway affects, directly or indirectly, the flow of metabolites through all the others. Cells and Organisms Maintain a Dynamic SteadyState The pathways of glucose metabolism provide, in the catabolic direction, the energy essential to oppose the forces of entropy and, in the anabolic direction, biosynthetic precursors and a storage form of metabolic energy. These reactions are so important to survival that very complex regulatory mechanisms have evolved to ensure that metabolites move through each pathway in the correct direction and at the correct rate to match exactly the cell’s or the organism’s changing circumstances. By a variety of mechanisms operating on different time scales, adjustments are made in the rate of metabolite flow through an entire pathway when external circumstances change. Circumstances do change, sometimes dramatically. The availability of oxygen may decrease due to hypoxia (diminished delivery of oxygen to tissues) or ischemia (diminished flow of blood to tissues). Wound healing requires huge amounts of energy and biosynthetic precursors. The relative proportions of carbohydrate, fat, and protein in the diet vary from meal to meal, and the supply of fuels obtained in the diet is intermittent, requiring metabolic adjustments between meals and during periods of starvation. Fuels such as glucose enter a cell, and waste products such as CO2 leave, but the mass and the gross composition of a typical cell, organ, or adult animal do not change appreciably over time; cells and organisms exist in a dynamic steady state. For each metabolic reaction in a pathway, the substrate is provided by the preceding reaction at the same rate at which it is converted to product. Thus, although the rate (v) of metabolite flow, or flux, through this step of the pathway may be high and variable, the concentration of substrate, S, remains constant. So, for the two-step reaction A v1→ S v2→ P when v1=v2, [S] is constant. For example, changes in v1 for the entry of glucose from various sources into the blood are balanced by changes in v2 for the uptake of glucose from the blood into various tissues, so the concentration of glucose in the blood ([S]) is held nearly constant at 5 mM. This is homeostasis for blood glucose. The failure of homeostatic mechanisms is o en at the root of human disease. In diabetes mellitus, for example, the regulation of blood glucose concentration is defective as a result of the lack of or insensitivity to insulin, with profound medical consequences. In the course of evolution, organisms have acquired a remarkable collection of regulatory mechanisms for maintaining homeostasis at the molecular, cellular, and organismal levels, as reflected in the proportion of genes that encode regulatory machinery. In humans, about 2,500 genes (~12% of all genes) encode regulatory proteins, including a variety of receptors, regulators of gene expression, and more than 500 different protein kinases! In many cases, the regulatory mechanisms overlap: one enzyme is subject to regulation by several different mechanisms. Both the Amount and the Catalytic Activity of anEnzyme Can Be Regulated The flux through an enzyme-catalyzed reaction can be modulated by changes in the number of enzyme molecules or by changes in the catalytic activity of each enzyme molecule already present. Such changes occur on time scales from milliseconds to many hours, in response to signals from within or outside the cell. Very rapid allosteric changes in enzyme activity are generally triggered locally, by changes in the local concentration of a small molecule — a substrate of the pathway in which that reaction is a step (say, glucose for glycolysis), a product of the pathway (ATP from glycolysis), or a key metabolite or cofactor (such as NADH) that indicates the cell’s metabolic state. Second messengers (such as cyclic AMP and Ca2+) generated intracellularly in response to extracellular signals (hormones, cytokines, and so forth) also mediate allosteric regulation, on a slightly slower time scale set by the rate of the signal-transduction mechanism (see Chapter 12). Extracellular signals (Fig. 13-29, ) may be hormonal (insulin or epinephrine, for example) or neuronal (acetylcholine), or may be growth factors or cytokines. The number of molecules of a given enzyme in a cell is a function of the relative rates of synthesis and degradation of that enzyme. The rate of synthesis can be adjusted by the activation (in response to some outside signal) of a transcription factor (Fig. 13- 29, ; described in more detail in Chapter 28). Transcription factors are nuclear proteins that, when activated, bind specific DNA regions (response elements) near a gene’s promoter (its transcriptional starting point) and activate or repress the transcription of that gene, leading to increased or decreased synthesis of the encoded protein. Activation of a transcription factor is sometimes the result of its binding of a specific ligand and sometimes the result of its phosphorylation or dephosphorylation. Each gene is controlled by one or more response elements that are recognized by specific transcription factors. Genes that have several response elements are therefore controlled by several different transcription factors responding to several different signals. Groups of genes encoding proteins that act together, such as the enzymes of glycolysis, o en share common response element sequences, so that a single signal, acting through a particular transcription factor, turns all of these genes on and off together. FIGURE 13-29 Factors affecting the activity of enzymes. The total activity of an enzyme can be changed by altering the number of its molecules in the cell, or its effective activity in a subcellular compartment ( through ), or by modulating the activity of existing molecules ( through ), as detailed in the text. An enzyme may be influenced by a combination of such factors. The stability of messenger RNAs — their resistance to degradation by cellular ribonucleases (Fig. 13-29, ) — varies, and the amount of a given mRNA in the cell is a function of its rates of synthesis and degradation (Chapter 26). The rate at which an mRNA is translated into a protein by ribosomes (Fig. 13-29, ) is also regulated, and depends on several factors described in detail in Chapter 27. Note that an n-fold increase in an mRNA does not always mean an n-fold increase in its protein product. Once synthesized, protein molecules have a finite lifetime, which may range from minutes to many days (Table 13-8). The rate of protein degradation (Fig. 13-29, ) differs from one protein to another and depends on the conditions in the cell. Some proteins are tagged by the covalent attachment of ubiquitin for degradation in proteasomes, as discussed in Chapter 27 (see, for example, the case of cyclin, in Fig. 12-38). Rapid turnover (synthesis followed by degradation) is energetically expensive, but proteins with a short half-life can reach new steady-state levels much faster than those with a long half-life, and the benefit of this quick responsiveness must balance or outweigh the cost to the cell. TABLE 13-8 Average Half-Life of Proteins in Mammalian Tissues Tissue Average half-life (days) Liver 0.9 Kidney 1.7 Heart 4.1 Brain 4.6 Muscle 10.7 Yet another way to alter the effective activity of an enzyme is to sequester the enzyme and its substrate in different compartments (Fig. 13-29, ). In muscle, for example, hexokinase cannot act on glucose until the sugar enters the myocyte from the blood, and the rate at which it enters depends on the activity of glucose transporters (see Table 11-1) in the plasma membrane. Within cells, membrane-bounded compartments segregate certain enzymes and enzyme systems, and the transport of substrate across these intracellular membranes may be the limiting factor in enzyme action. By these several mechanisms for regulating enzyme level, cells can dramatically change their complement of enzymes in response to changes in metabolic circumstances. In vertebrates, liver is the most adaptable tissue; a change from a high-carbohydrate diet to a high-lipid diet, for example, affects the transcription of hundreds of genes and thus the levels of hundreds of proteins. These global changes in gene expression can be quantified in the entire complement of mRNAs (the transcriptome) or protein (proteome) of a cell type or organ, offering great insights into metabolic regulation. The effect of changes in the proteome is o en a change in the total ensemble of low molecular weight metabolites, the metabolome (Fig. 13- 30). The metabolome of E. coli growing on glucose is dominated by a few classes of metabolites: glutamate (49%); nucleotides (mainly ribonucleoside triphosphates) (15%); intermediates of glycolysis, the citric acid cycle, and the pentose phosphate pathway (central pathways of carbon metabolism) (15%); and redox cofactors and glutathione (9%). FIGURE 13-30 The metabolome of E. coli growing on glucose. Summary of the relative molar abundance of 103 metabolites as measured by a combination of liquid chromatography and tandem mass spectrometry (LC-MS/MS). For reference, the absolute concentration of glutamate in living cells is 9.6 m . [Data from B. D. Bennett et al., Nature Chem. Biol. 5:593, 2009, Fig. 1.] Once the regulatory mechanisms that involve protein synthesis and degradation have produced a certain number of molecules of each enzyme in a cell, the activity of those enzymes can be further regulated in several other ways: by the concentration of substrate, the presence of allosteric effectors, covalent modifications, or binding of regulatory proteins — all of which can change the activity of an individual enzyme molecule (Fig. 13-29, to ). All enzymes are sensitive to the concentration of their substrate(s) (Fig. 13-29, ). Recall that in the simplest case (an enzyme that follows Michaelis-Menten kinetics), the initial rate of the reaction is half-maximal when the substrate is present at a concentration equal to Km (that is, when the enzyme is half-saturated with substrate). Activity drops off at lower [S], and when [S] ≪ Km, the reaction rate is linearly dependent on [S]. The relationship between [S] and Km is important because intracellular concentrations of substrate are o en in the same range as, or lower than, Km. The activity of hexokinase, for example, changes with [glucose], and intracellular [glucose] varies with the concentration of glucose in the blood. As we will see, the different forms (isozymes) of hexokinase have different Km values and are therefore affected differently by changes in intracellular [glucose], in ways that make sense physiologically. For a number of phosphoryl transfers from ATP, and for redox reactions using NADPH or NAD+, the metabolite concentration is well above the Km (Fig. 13-31); these cofactors are not likely to be limiting in such reactions. FIGURE 13-31 Comparison of Km and substrate concentration for some metabolic enzymes. Measured metabolite concentrations for E. coli growing on glucose are plotted against the known Km for enzymes that consume that metabolite. The solid line is the line of unity (where metabolite concentration =Km), and the dashed lines each denote a 10-fold deviation from the line of unity. [Data from B. D. Bennett et al., Nature Chem. Biol. 5:593, 2009, Fig. 2.] Enzyme activity can be either increased or decreased by an allosteric effector (Fig. 13-29, ; see Fig. 6-37). Allosteric effectors typically convert hyperbolic kinetics to sigmoid kinetics, or vice versa (see Fig. 14-24b, for example). In the steepest part of the sigmoid curve, a small change in the concentration of substrate, or of allosteric effector, can have a large impact on reaction rate. Recall from Chapter 5 (p. 157) that the cooperativity of an allosteric protein can be expressed as a Hill coefficient, with higher coefficients meaning greater cooperativity. For an allosteric enzyme with a Hill coefficient of 4, activity increases from 10% Vmax to 90% Vmax with only a 3-fold increase in [S], compared with the 81-fold rise in [S] needed by an enzyme with no cooperative effects (Hill coefficient of 1; Table 13-9). TABLE 13-9 Relationship between Hill Coefficient and the Effect of Substrate Concentration on Reaction Rate for Allosteric Enzymes Hill coefficient (nH) Required change in [S] to increase V0 from 10% to 90% Vmax 0.5 ×6,600 1.0 ×81 2.0 ×9 3.0 ×4.3 4.0 ×3 Covalent modifications of enzymes or other proteins (Fig. 13-29, ) occur within seconds or minutes of a regulatory signal, typically an extracellular signal. By far the most common modifications are phosphorylation and dephosphorylation (Fig. 13- 32); up to half the proteins in a eukaryotic cell are phosphorylated under some circumstances. Phosphorylation by a specific protein kinase may alter the electrostatic features of an enzyme’s active site, cause movement of an inhibitory region of the enzyme out of the active site, alter the enzyme’s interaction with other proteins, or force conformational changes that translate into changes in Vmax or Km. For covalent modification to be useful in regulation, the cell must be able to restore the altered enzyme to its original activity state. A family of phosphoprotein phosphatases, at least some of which are themselves under regulation, catalyzes the dephosphorylation of proteins. FIGURE 13-32 Protein phosphorylation and dephosphorylation. Protein kinases transfer a phosphoryl group from ATP to a Ser, Thr, or Tyr residue in an enzyme or other protein substrate. Protein phosphatases remove the phosphoryl group as Pi. Finally, many enzymes are regulated by association with and dissociation from another, regulatory protein (Fig. 13-29, ). For example, the cyclic AMP–dependent protein kinase (PKA; see Fig. 12-6) is inactive until cAMP binding separates catalytic from regulatory (inhibitory) subunits of the enzyme. These several mechanisms for altering the flux through a step in a metabolic pathway are not mutually exclusive. It is very common for a single enzyme to be regulated at the level of transcription and by both allosteric and covalent mechanisms. The combination provides fast, smooth, effective regulation in response to a very wide array of perturbations and signals. In the discussions that follow, it is useful to think of changes in enzymatic activity as serving two distinct though complementary roles. We use the term metabolic regulation to refer to processes that serve to maintain homeostasis at the molecular level — to hold some cellular parameter (concentration of a metabolite, for example) at a steady level over time, even as the flow of metabolites through the pathway changes. The term metabolic control refers to a process that leads to a change in the output of a metabolic pathway over time, in response to some outside signal or change in circumstances. The distinction, although useful, is not always easy to make. Reactions Far from Equilibrium in Cells AreCommon Points of Regulation For some steps in a metabolic pathway the reaction is close to equilibrium, with the cell in its dynamic steady state (Fig. 13-33). The net flow of metabolites through these steps is the small difference between the rates of the forward and reverse reactions, rates that are very similar when a reaction is near equilibrium. Small changes in substrate or product concentration can produce large changes in the net rate, and they can even change the direction of the net flow. We can identify these near- equilibrium reactions in a cell by comparing the mass-action ratio, Q, with the equilibrium constant for the reaction, K′eq. Recall that for the reaction A+B→ C+D, Q=[C][D]/[A][B]. In practice, when Q and K′eq are within 1 to 2 orders of magnitude of each other, the reaction is near equilibrium. This is the case for more than half of the enzymes in the glycolytic pathway, for example (Table 13- 10). FIGURE 13-33 Near-equilibrium and nonequilibrium steps in a metabolic pathway. Steps and of this pathway are near equilibrium in the cell; for each step, the rate (V) of the forward reaction is only slightly greater than the reverse rate, so the net forward rate (10) is relatively low and the free-energy change, ∆G, is close to zero. An increase in [C] or [D] can reverse the direction of these steps. Step is maintained in the cell far from equilibrium; its forward rate greatly exceeds its reverse rate. The net rate of step (10) is much larger than the reverse rate (0.01) and is identical to the net rates of steps and when the pathway is operating in the steady state. Step has a large, negative ∆G. TABLE 13-10 Equilibrium Constants, Mass-Action Ratios, and Free-Energy Changes for Enzymes of Carbohydrate Metabolism Enzyme K′eq Mass-action ratio, Q Reaction near equilibrium in vivo? ΔG′° (kJ/mol) Δ G (kJ/mol) in heartLiver Heart Hexokinase 1×103 2×10−2 8×10−2 No −17 −27 PFK-1 1.0×103 9×10−2 3×10−2 No −14 −23 Aldolase 1.0×10−4 1.2×10−6 9×10−6 Yes +24 −6.0 Triose phosphate isomerase 4×10−2 — 2.4×10−1 Yes +7.5 +3.8 Glyceraldehyde 3- phosphate dehydrogenase + phosphoglycerate kinase 2×103 6×102 9.0 Yes −13 +3.5 Phosphoglycerate mutase 1×10−1 1×10−1 1.2×10−1 Yes +4.4 +0.6 Enolase 3 2.9 1.4 Yes −3.2 −0.5 Pyruvate kinase 2×104 7×10−1 40 No −31 −17 Phosphohexose isomerase 4×10−1 3.1×10−1 2.4×10−1 Yes +2.2 −1.4 Pyruvate carboxylase + PEP carboxykinase 7 1×10−3 — No −5.0 −23 Glucose 6- phosphatase 8.5×102 1.2×102 — Yes −17 −5.0 Data for K′eq and Q from E. A. Newsholme and C. Start, Regulation in Metabolism, pp. 97, 263, Wiley Press, 1973. ∆G and ΔG′° were calculated from these data. For simplicity, any reaction for which the absolute value of the calculated ∆G is less than 6 is considered near equilibrium. Data not available. a b b b a b Other reactions are far from equilibrium in the cell. For example, K′eq for the phosphofructokinase-1 (PFK-1) reaction is about 1,000, but Q ([fructose 1,6- bisphosphate][ADP]/[fructose 6-phosphate][ATP]) in a hepatocyte in the steady state is about 0.1 (Table 13-10). It is because the reaction is so far from equilibrium in the cell that the process is exergonic under cellular conditions and tends to go in the forward direction. The reaction is held far from equilibrium because, under prevailing cellular conditions of substrate, product, and effector concentrations, the rate of conversion of fructose 6-phosphate to fructose 1,6-bisphosphate is limited by the activity of PFK-1. PFK-1 activity itself is limited by the number of PFK-1 molecules present and by the actions of allosteric effectors. Thus, the net forward rate of the enzyme-catalyzed reaction is equal to the net flow of glycolytic intermediates through other steps in the pathway, and the reverse flow through PFK- 1 remains near zero. The cell cannot allow reactions with large equilibrium constants to reach equilibrium. If [fructose 6-phosphate], [ATP], and [ADP] in the cell were held at typical levels (low millimolar concentrations) and the PFK-1 reaction were allowed to reach equilibrium by an increase in [fructose 1,6-bisphosphate], the concentration of fructose 1,6-bisphosphate would rise into the molar range, wreaking osmotic havoc on the cell. Consider another case: if the reaction ATP→ ADP+Pi were allowed to approach equilibrium in the cell, the actual free-energy change (ΔG) for that reaction (ΔGp; see Worked Example 13-2, p. 480) would approach zero, and ATP would lose the high phosphoryl group transfer potential that makes it valuable to the cell. It is therefore essential that enzymes catalyzing ATP breakdown and other highly exergonic reactions in a cell be sensitive to regulation, so that when metabolic changes are forced by external circumstances, the flow through these enzymes will be adjusted to ensure that [ATP] remains far above its equilibrium level. When such metabolic changes occur, the activities of enzymes in all interconnected pathways adjust to keep these critical steps away from equilibrium. Thus, not surprisingly, many enzymes that catalyze highly exergonic reactions are subject to a variety of subtle regulatory mechanisms. The multiplicity of these adjustments is so great that we cannot predict by examining the properties of any one enzyme in a pathway whether that enzyme has a strong influence on net flow through the entire pathway. Adenine Nucleotides Play Special Roles inMetabolic Regulation A er the protection of its DNA from damage, perhaps nothing is more important to a cell than maintaining a constant supply and concentration of ATP. Many ATP-using enzymes have Km values between 0.1 and 1 mM, and the ATP concentration in a typical cell is about 5 to 10 mM (Fig. 13-31). If [ATP] were to drop significantly, these enzymes would be less than fully saturated by their substrate (ATP), and the rates of hundreds of reactions that involve ATP would decrease (Fig. 13-34); the cell would probably not survive this kinetic effect on so many reactions. FIGURE 13-34 Effect of ATP concentration on the initial reaction velocity of a typical ATP-dependent enzyme. These experimental data yield a Km for ATP of 5 m . The concentration of ATP in animal tissues is ~5 m . There is also an important thermodynamic effect of lowered [ATP]. Because ATP is converted to ADP or AMP when “spent” to accomplish cellular work, the [ATP]/[ADP] ratio profoundly affects all reactions that employ these cofactors. The same is true for other important cofactors, such as NADH/NAD+ and NADPH/NADP+. For example, consider the reaction catalyzed by hexokinase: ATP+glucose→ ADP+glucose 6-phosphate K′eq= =2×103 Note that this expression holds true only when reactants and products are at their equilibrium concentrations, where ΔG=0. At any other set of concentrations, ΔG is not zero. Recall (from Section 13.1) that the ratio of products to substrates (the mass- action ratio, Q) determines the magnitude and sign of ΔG and therefore the driving force, ΔG, of the reaction: ΔG=ΔG′°+RT ln Because an alteration of this driving force profoundly influences every reaction that involves ATP, organisms have evolved under strong pressure to develop regulatory mechanisms responsive to the [ATP]/[ADP] ratio. AMP concentration is an even more sensitive indicator of a cell’s energetic state than is [ATP]. Normally, cells have a far higher concentration of ATP (5 to 10 mM) than of AMP (<0.1 mM). When some process (say, muscle contraction) consumes ATP, AMP is produced in two steps. First, hydrolysis of ATP produces ADP, then the reaction catalyzed by adenylate kinase produces AMP: 2ADP→ AM P+ATP If ATP is consumed such that its concentration drops 10%, the relative increase in [AMP] is much greater than that of [ADP] (Table 13-11). It is not surprising, therefore, [ADP]eq[glucose 6-phosphate]eq [ATP]eq[glucose]eq [ADP][glucose 6-phosphate] [ATP][glucose] that many regulatory processes are keyed to changes in [AMP]. Probably the most important mediator of regulation by AMP is AMP-activated protein kinase (AMPK), which responds to an increase in [AMP] by phosphorylating key proteins and thus regulating their activities. (AMPK is not to be confused with the cyclic AMP– dependent protein kinase PKA; see Section 12.2.) The rise in [AMP] may be caused by a reduced nutrient supply or by increased physical exercise. AMP activates AMPK allosterically, which increases glucose transport and activates glycolysis and fatty acid oxidation, while suppressing energy-requiring processes such as the synthesis of glycogen, fatty acids, cholesterol, and protein. All of the changes effected by AMPK serve to raise [ATP] and lower [AMP]. In Chapter 23, we discuss the role of AMPK in balancing anabolism and catabolism in the whole organism. TABLE 13-11 Relative Changes in [ATP] and [AMP] When ATP Is Consumed Adenine nucleotide Concentration before ATP depletion (m ) Concentration a er ATP depletion (m ) Relative change ATP 5.0 4.5 10% ADP 1.0 1.0 0 AMP 0.1 0.6 600% SUMMARY 13.5 Regulation of MetabolicPathways In a metabolically active cell in a steady state, intermediates are formed and consumed at equal rates. When a transient perturbation alters the rate of formation or consumption of a metabolite, compensating changes in enzyme activities return the system to the steady state. Cells regulate their metabolism by a variety of mechanisms over a time scale ranging from less than a millisecond to days, either by changing the activity of existing enzyme molecules or by changing the number of molecules of a specific enzyme. Various signals activate or inactivate transcription factors, which act in the nucleus to regulate gene expression. Changes in the transcriptome lead to changes in the proteome, and ultimately in the metabolome of a cell or tissue. In multistep processes such as glycolysis, certain reactions are essentially at equilibrium in the steady state; the rates of these reactions rise and fall with substrate concentration. Other reactions are far from equilibrium; these steps are typically the points of regulation of the overall pathway. Regulatory mechanisms maintain nearly constant levels of key metabolites such as ATP and NADH in cells and glucose in the blood, while matching the use or production of glucose to the organism’s changing needs. The levels of ATP and AMP are a sensitive reflection of a cell’s energy status, and when the [ATP]/[AMP] ratio decreases, the AMP-activated protein kinase (AMPK) triggers a variety of cellular responses to raise [ATP] and lower [AMP]. Chapter Review KEY TERMS Terms in bold are defined in the glossary. autotroph heterotroph metabolite intermediary metabolism catabolism anabolism energy transduction free energy, G exergonic endergonic enthalpy, H exothermic endothermic entropy, S standard transformed constants mass-action ratio, Q homolytic cleavage radical heterolytic cleavage nucleophile electrophile carbanion carbocation aldol condensation Claisen condensation kinases phosphorylation potential (ΔGp) thioester adenylylation inorganic pyrophosphatase nucleoside diphosphate kinase adenylate kinase creatine kinase phosphagens electromotive force (emf) conjugate redox pair dehydrogenases reducing equivalent standard reduction potential (E°) pyridine nucleotide oxidoreductase flavoprotein flavin nucleotides glucose 6-phosphate homeostasis transcription factor response element turnover transcriptome proteome metabolome metabolic regulation metabolic control adenylate kinase AMP-activated protein kinase (AMPK) PROBLEMS 1. Entropy Changes during Egg Development Consider a system consisting of an egg in an incubator. The white and yolk of the egg contain proteins, carbohydrates, and lipids. If fertilized, the egg transforms from a single cell to a complex organism. Discuss this irreversible process in terms of the entropy changes in the system and surroundings. Be sure that you first clearly define the system and surroundings. 2. Calculation of ΔG′° from an Equilibrium Constant Calculate the standard free-energy change for each of the three metabolically important enzyme-catalyzed reactions, using the equilibrium constants given for the reactions at 25 °C and pH 7.0. a. b. c. 3. Calculation of the Equilibrium Constant from ΔG′° Calculate the equilibrium constant K′eq for each of the three reactions at pH 7.0 and 25 °C, using the ΔG′° values in Table 13-4. a. b. c. 4. Experimental Determination of K′eq and ΔG′° Incubating a 0.1 M solution of glucose 1-phosphate at 25 °C with a catalytic amount of phosphoglucomutase transforms some of the glucose 1-phosphate to glucose 6-phosphate. At equilibrium, the concentrations of the reaction components are G lucose 1-phosphate ⇌ glucose 6-phosphate 4.5× 10−3M 9.6× 10−2M Calculate K′eq and ΔG′° for this reaction. 5. Experimental Determination of ΔG′° for ATP Hydrolysis A direct measurement of the standard free-energy change associated with the hydrolysis of ATP is technically demanding because the minute amount of ATP remaining at equilibrium is difficult to measure accurately. The value of ΔG′° can be calculated indirectly, however, from the equilibrium constants of two other enzymatic reactions having less favorable equilibrium constants: G lucose 6-phosphate+ H2O → glucose+ Pi K′eq = 270 AT P + glucose→ AD P + glucose 6-phosphate K′eq = 890 Using this information for equilibrium constants determined at 25 °C, calculate the standard free energy of hydrolysis of ATP. 6. Difference between ΔG′° and Δ G Consider the interconversion shown, which occurs in glycolysis (Chapter 14): Fructose 6-phosphate ⇌ glucose 6-phosphate K′eq = 1.97 a. What is ΔG′° for the reaction (K′eq measured at 25 °C)? b. If the concentration of fructose 6-phosphate is adjusted to 1.5 M and that of glucose 6-phosphate is adjusted to 0.50 M, what is ΔG? c. Why are ΔG′° and ΔG different? 7. Free Energy of Hydrolysis of CTP Compare the structure of the nucleoside triphosphate CTP with the structure of ATP. Now predict the K′eq and ΔG′° for the reaction: AT P + CD P → AD P + CT P 8. Dependence of Δ G on pH The free energy released by the hydrolysis of ATP under standard conditions is −30.5 kJ/mol. If ATP is hydrolyzed under standard conditions except at pH 5.0, is more or less free energy released? Explain. 9. The ΔG′° for Coupled Reactions Glucose 1-phosphate is converted into fructose 6-phosphate in two successive reactions: G lucose 1-phosphate→ glucose 6-phosphate G lucose 6-phosphate→ fructose 6-phosphate Using the ΔG′° values in Table 13-4, calculate the equilibrium constant, K′eq, for the sum of the two reactions: G lucose 1-phosphate → fructose 6-phosphate 10. Effect of [ATP]/[ADP] Ratio on Free Energy of Hydrolysis of ATP Using Equation 13-4, plot ΔG against ln Q (mass-action ratio) at 25 °C for the concentrations of ATP, ADP, and Pi in the table shown. ΔG′° for the reaction is −30.5 kJ/mol. Use the resulting plot to explain why metabolism is regulated to keep the ratio [ATP]/[ADP] high. Concentration (m ) ATP 5 3 1 0.2 5 ADP 0.2 2.2 4.2 5.0 25 Pi 10 12.1 14.1 14.9 10 11. Strategy for Overcoming an Unfavorable Reaction: ATP- Dependent Chemical Coupling The phosphorylation of glucose to glucose 6-phosphate is the initial step in the catabolism of glucose. The direct phosphorylation of glucose by Pi is described by the equation a. Calculate the equilibrium constant for this reaction at 37 °C. In the rat hepatocyte, the physiological concentrations of glucose and Pi are maintained at approximately 4.8 mM. What is the equilibrium concentration of glucose 6-phosphate obtained by the direct phosphorylation of glucose by Pi? Does this reaction represent a reasonable metabolic step for the catabolism of glucose? Explain. b. In principle at least, one way to increase the concentration of glucose 6-phosphate is to drive the equilibrium reaction to the right by increasing the intracellular concentrations of glucose and Pi. Assuming a fixed concentration of Pi at 4.8 mM, how high would the intracellular concentration of glucose have to be to give an equilibrium concentration of glucose 6-phosphate of 250 μ M (the normal physiological concentration)? Would this route be physiologically reasonable, given that the maximum solubility of glucose is less than 1 M? c. The phosphorylation of glucose in the cell is coupled to the hydrolysis of ATP; that is, part of the free energy of ATP hydrolysis is used to phosphorylate glucose: G lucose+ Pi → glucose 6-phosphate+ H2O ΔG′°= 13.8 kJ/m Calculate K′eq at 37 °C for the overall reaction. For the ATP-dependent phosphorylation of glucose, what concentration of glucose is needed to achieve a 250 μ M intracellular concentration of glucose 6-phosphate when the concentrations of ATP and ADP are 3.38 mM and 1.32 mM, respectively? Does this coupling process provide a feasible route, at least in principle, for the phosphorylation of glucose in the cell? Explain. d. Although coupling ATP hydrolysis to glucose phosphorylation makes thermodynamic sense, we have not yet specified how this coupling is to take place. Given that coupling requires a common intermediate, one conceivable route is to use ATP hydrolysis to raise the intracellular concentration of Pi and thus drive the unfavorable phosphorylation of glucose by Pi. Is this a reasonable route? (Think about the solubility product, Ksp, of metabolic intermediates.) e. The ATP-coupled phosphorylation of glucose is catalyzed in hepatocytes by the enzyme glucokinase. This enzyme binds ATP and glucose to form a glucose-ATP-enzyme complex, and the phosphoryl group is transferred directly from ATP to glucose. Explain the advantages of this route. (1) G lucose+ Pi→ glucose 6-phosphate+ H2O ΔG′°= 13.8 kJ/mo (2) AT P + H2O → AD P + Pi ΔG′°= −30.5 kJ––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– Sum : G lucose+ AT P → glucose 6-phosphate+ AD 12. Calculations of ΔG′° for ATP-Coupled Reactions From data in Table 13-6, calculate the ΔG′° value for each reaction: a. Phosphocreatine+ AD P → creatine+ AT P b. AT P + fructose→ AD P + fructose 6-phosphate 13. Coupling ATP Cleavage to an Unfavorable Reaction To explore the consequences of coupling ATP hydrolysis under physiological conditions to a thermodynamically unfavorable biochemical reaction, consider the hypothetical transformation X → Y, for which ΔG′°= 20.0 kJ/mol. a. What is the ratio [Y]/[X] at equilibrium? b. Suppose X and Y participate in a sequence of reactions during which ATP is hydrolyzed to ADP and Pi. The overall reaction is X + AT P + H2O → Y + AD P + Pi Calculate [Y]/[X] for this reaction at equilibrium. Assume that the temperature is 25.0 °C and the equilibrium concentrations of ATP, ADP, and Pi are 1 M. c. We know that [ATP], [ADP], and [Pi] are not 1 M under physiological conditions. Calculate [Y]/[X] for the ATP- coupled reaction when the values of [ATP], [ADP], and [Pi] are those found in rat myocytes (Table 13-5). 14. Calculations of Δ G at Physiological Concentrations Calculate the actual, physiological ΔG for the reaction Phosphocreatine+ AD P → creatine+ AT P at 37 °C, as it occurs in the cytosol of neurons, with phosphocreatine at 4.7 mM, creatine at 1.0 mM, ADP at 0.73 mM, and ATP at 2.6 mM. 15. Free Energy Required for ATP Synthesis under Physiological Conditions In the cytosol of rat hepatocytes, the temperature is 37 °C and the mass-action ratio, Q, is = 5.33× 102 M −1 Calculate the free energy required to synthesize ATP in a rat hepatocyte. 16. Chemical Logic In the glycolytic pathway, a six-carbon sugar (fructose 1,6-bisphosphate) is cleaved to form two three-carbon sugars, which undergo further metabolism. In this pathway, an isomerization of glucose 6-phosphate to fructose 6-phosphate (as shown in the diagram) occurs two steps before the cleavage reaction. The intervening step is phosphorylation of fructose 6- phosphate to fructose 1,6-bisphosphate (p. 516). [AT P] [AD P][Pi] What does the isomerization step accomplish from a chemical perspective? (Hint: Consider what might happen if the C—C bond cleavage were to proceed without the preceding isomerization.) 17. Enzymatic Reaction Mechanisms I Lactate dehydrogenase is one of the many enzymes that require NADH as coenzyme. It catalyzes the conversion of pyruvate to lactate:
Draw the mechanism of this reaction (show electron-pushing arrows). (Hint: This is a common reaction throughout metabolism; the mechanism is similar to that catalyzed by other dehydrogenases that use NADH, such as alcohol dehydrogenase.) 18. Enzymatic Reaction Mechanisms II Biochemical reactions o en look more complex than they really are. In the pentose phosphate pathway (Chapter 14), sedoheptulose 7-phosphate and glyceraldehyde 3-phosphate react to form erythrose 4- phosphate and fructose 6-phosphate in a reaction catalyzed by transaldolase. Draw a mechanism for this reaction (show electron-pushing arrows). (Hint: Take another look at aldol condensations, then consider the name of this enzyme.) 19. Recognizing Reaction Types For each pair of biomolecules, identify the type of reaction (oxidation-reduction, hydrolysis, isomerization, group transfer, or internal rearrangement) required to convert the first molecule to the second. In each case, indicate the general type of enzyme and cofactor(s) or reactants that would be required, and any other products that would result. a. b.
c. d.
e. f. g.
20. Effect of Structure on Group Transfer Potential Some invertebrates contain phosphoarginine. Is the standard free energy of hydrolysis of this molecule more similar to that of glucose 6-phosphate or of ATP? Explain your answer. 21. Polyphosphate as a Possible Energy Source The standard free energy of hydrolysis of inorganic polyphosphate (polyP) is about −20 kJ/mol for each Pi released. We calculated in Worked Example 13-2 that, in a cell, it takes about 50 kJ/mol of energy to synthesize ATP from ADP and Pi. Is it feasible for a cell to use polyphosphate to synthesize ATP from ADP? Explain your answer. 22. Daily ATP Utilization by Human Adults a. The synthesis of ATP from ADP and Pi requires a total of 30.5 kJ/mol of free energy when the reactants and products are at 1 M concentrations and the temperature is 25 °C (standard state). However, the actual physiological concentrations of ATP, ADP, and Pi are not 1 M, and the physiological temperature is 37 °C. Thus, the free energy required to synthesize ATP under physiological conditions is different from ΔG′°. Calculate the free energy required to synthesize ATP in the human hepatocyte when the physiological concentrations of ATP, ADP, and Pi are 3.5, 1.50, and 5.0 mM, respectively. b. A 68 kg (150 lb) adult requires a caloric intake of 2,000 kcal (8,360 kJ) of food per day (24 hours). The body metabolizes the food and uses the free energy to synthesize ATP, which then provides energy for the body’s daily chemical and mechanical work. Assuming that the efficiency of converting food energy into ATP is 50%, calculate the weight of ATP used by a human adult in 24 hours. What percentage of the body weight does this represent? c. Although adults synthesize large amounts of ATP daily, their body weight, structure, and composition do not change significantly during this period. Explain this apparent contradiction. 23. Rates of Turnover of γ and β Phosphates of ATP A er adding a small amount of ATP labeled with radioactive phosphorus in the terminal position, [γ-32P]ATP, to a yeast extract, a researcher finds about half of the 32P activity in Pi within a few minutes, but the concentration of ATP remains unchanged. Explain. She then carries out the same experiment using ATP labeled with 32P in the central position, [β-32P]ATP, but the 32P does not appear in Pi within such a short time. Why? 24. Cleavage of ATP to AMP and PPi during Metabolism Synthesis of the activated form of acetate (acetyl-CoA) is carried out in an ATP-dependent process: Acetate+ CoA + AT P → acetyl-CoA + AM P + PPi a. The ΔG′° for hydrolysis of acetyl-CoA to acetate and CoA is −32.2 kJ/mol. The ΔG′° for hydrolysis of ATP to AMP and PPi is −30.5 kJ/mol. Calculate ΔG′° for the ATP- dependent synthesis of acetyl-CoA. b. Almost all cells contain the enzyme inorganic pyrophosphatase, which catalyzes the hydrolysis of PPi to Pi. What effect does the presence of this enzyme have on the synthesis of acetyl-CoA? Explain. 25. Activation of a Fatty Acid by Reaction with Coenzyme A In the reaction sequence for fatty acid breakdown, coenzyme A (CoA), with its thiol (— SH) group, joins to the fatty acid as a thiol ester, as ATP is converted into AMP and PPi: R—COO− + AT P + CoA—SH → AM P + PPi+ R—CO—S—CoA The oxidation of fatty acids as fuels requires two steps. The first step transfers an activating group from ATP to the carboxyl group of the fatty acid. In the second step, CoA—SH displaces the activating group to form fatty acyl-S—CoA. Given the known products of the reaction, what is the activating group? 26. Energy for H+ Pumping The parietal cells of the stomach lining contain membrane “pumps” that transport hydrogen ions from the cytosol (pH 7.0) into the stomach, contributing to the acidity of gastric juice (pH 1.0). Calculate the free energy required to transport 1 mol of hydrogen ions through these pumps. (Hint: See Chapter 11.) Assume a temperature of 37 °C. 27. Most-Reduced Carbon Compounds Arrange the four structures in order from most reduced to most oxidized. a. R—CH2—CH2—OH b. R—CH2—COO− c. R—CH2—CHO d. R—CH2—CH3 28. Standard Reduction Potentials The standard reduction potential, E ′°, of any redox pair is defined for the half-cell reaction Oxidizing agent+ n electrons → reducing agent The E ′° values for the NAD +/NAD H and pyruvate/lactate conjugate redox pairs are −0.32 V and −0.19 V, respectively. a. Which redox pair has the greater tendency to lose electrons? Explain. b. Which pair is the stronger oxidizing agent? Explain. c. Beginning with 1 M concentrations of each reactant and product at pH 7 and 25 °C, in which direction will the following reaction proceed? Pyruvate+ NAD H + H+ ⇌ lactate+ NAD + d. What is the standard free-energy change (ΔG′°) for the conversion of pyruvate to lactate? e. What is the equilibrium constant (K′eq) for this reaction? 29. Simple Biobattery Suppose you set up a simple battery using half-reactions as pictured in Figure 13-23. One electrode contains pyruvate and lactate at 1 mM, and the other electrode contains fumarate and succinate at 1 mM (see Table 13-7). a. In which direction will electrons initially flow? b. Calculate the standard reduction potential and standard free-energy change for your biological battery. c. When a flashlight battery “runs out,” net electron movement has essentially ended. What is the equivalent situation for your biobattery? 30. Energy Span of the Respiratory Chain Electron transfer in the mitochondrial respiratory chain may be represented by the net reaction equation NAD H + H+ + O2 ⇌ H2O + NAD + a. Calculate ΔE ′° for the net reaction of mitochondrial electron transfer. Use E ′° values in Table 13-7. b. Calculate ΔG′° for this reaction. c. How many ATP molecules can theoretically be generated by this reaction if the free energy of ATP synthesis under 1 2 cellular conditions is 52 kJ/mol? 31. Dependence of Electromotive Force on Concentrations Suppose that you place an electrode into solutions of various concentrations of NAD + and NADH at pH 7.0 and 25 °C. Calculate the electromotive force (in volts) registered by the electrode when immersed in each solution, with reference to a half-cell of E ′° 0.00 V. a. 1.0 mM NAD + and 10 mM NADH b. 1.0 mM NAD + and 1.0 mM NADH c. 10 mM NAD + and 1.0 mM NADH 32. Electron Affinity of Compounds List the four compounds or reactions in order of increasing tendency to accept electrons: a. α-ketoglutarate+ CO2 (yielding isocitrate) b. oxaloacetate c. O2 d. NAD P+ 33. Direction of Oxidation-Reduction Reactions Which of the reactions listed would you expect to proceed in the direction shown, under standard conditions, in the presence of the appropriate enzymes? a. M alate+ NAD + → oxaloacetate+ NAD H + H+ b. Acetoacetate+ NAD H + H+ → β-hydroxybutyrate+ NAD + c. Pyruvate+ NAD H + H+ → lactate+ NAD + d. Pyruvate+ β-hydroxybutyrate → lactate+ acetoacetate e. M alate+ pyruvate → oxaloacetate+ lactate f. Acetaldehyde+ succinate→ ethanol+ fumarate 34. Measurement of Intracellular Metabolite Concentrations Measuring the concentrations of metabolic intermediates in a living cell presents great experimental difficulties — usually, a cell must be destroyed before metabolite concentrations can be measured. Yet enzymes catalyze metabolic interconversions very rapidly, so a common problem associated with these types of measurements is that the findings reflect not the physiological concentrations of metabolites but the equilibrium concentrations. To prevent changes in metabolite concentrations during sample preparation, cells were quick- frozen in liquid nitrogen, then extracted under conditions that prevented enzymatic activity. The table gives the intracellular concentrations of the substrates and products of the phosphofructokinase-1 reaction in isolated rat heart tissue. Metabolite Concentration (μ M) Fructose 6-phosphate 87.0 Fructose 1,6-bisphosphate 22.0 ATP 11,400 ADP 1,320 Data from J. R. Williamson, J. Biol. Chem. 240:2308, 1965. Calculated as μ mol/mL of intracellular water. a. Calculate Q, [fructose 1,6-bisphosphate][ADP]/[fructose 6-phosphate][ATP], for the PFK-1 reaction under physiological conditions. b. Given a ΔG′° for the PFK-1 reaction of −14.2 kJ/mol, calculate the equilibrium constant for this reaction. a a c. Compare the values of Q and K′eq. Is the physiological reaction near or far from equilibrium? Explain. What does this experiment suggest about the role of PFK-1 as a regulatory enzyme? 35. Are All Metabolic Reactions at Equilibrium? a. Phosphoenolpyruvate (PEP) is one of the two phosphoryl group donors in the synthesis of ATP during glycolysis. In human erythrocytes, the steady-state concentration of ATP is 2.24 mM, that of ADP is 0.25 mM, and that of pyruvate is 0.051 mM. Calculate the concentration of PEP at 25 °C, assuming that the pyruvate kinase reaction (see Fig. 13-13) is at equilibrium in the cell. b. The physiological concentration of PEP in human erythrocytes is 0.023 mM. Compare this with the value obtained in (a). Explain the significance of this difference. 36. Michaelis Constant Km Compared with Substrate Concentration Malate synthase in E. coli catalyzes the reaction Acetyl-CoA + glyoxylate+ H2O → malate+ CoA-SH + H+ The experimentally measured Km for acetyl-CoA is 9× 106M . In a growing culture of E. coli, the measured concentration of acetyl-CoA is 6× 10−4M . Is malate synthase operating at its Vmax under these conditions? DATA ANALYSIS PROBLEM 37. Thermodynamics Can Be Tricky Thermodynamics is a challenging area of study and one with many opportunities for confusion. An interesting example is found in an article by Robinson, Hampson, Munro, and Vaney, published in Science in 1993. Robinson and colleagues studied the movement of small molecules between neighboring cells of the nervous system through cell-to-cell channels (gap junctions). They found that the dyes Lucifer yellow (a small, negatively charged molecule) and biocytin (a small zwitterionic molecule) moved in only one direction between two particular types of glia (nonneuronal cells of the nervous system). Dye injected into astrocytes would rapidly pass into adjacent astrocytes, oligodendrocytes, or Müller cells, but dye injected into oligodendrocytes or Müller cells passed slowly if at all into astrocytes. All of these cell types are connected by gap junctions. Although it was not a central point of their article, the authors presented a molecular model for how this unidirectional transport might occur, as shown in their Figure 3: The figure legend reads: “Model of the unidirectional diffusion of dye between coupled oligodendrocytes and astrocytes, based on differences in connection pore diameter. Like a fish in a fish trap, dye molecules (black circles) can pass from an astrocyte to an oligodendrocyte (A) but not back in the other direction (B).” Although this article clearly passed review at a well-respected journal, several letters to the editor (1994) followed, showing that Robinson and coauthors’ model violated the second law of thermodynamics. a. Explain how the model violates the second law. Hint: Consider what would happen to the entropy of the system if one started with equal concentrations of dye in the astrocyte and oligodendrocyte connected by the “fish trap” type of gap junctions. b. Explain why this model cannot work for small molecules, although it may allow one to catch fish. c. Explain why a fish trap does work for fish. d. Provide two plausible mechanisms for the unidirectional transport of dye molecules between the cells that do not violate the second law of thermodynamics. References Letters to the editor. 1994. Science 265:1017–1019. Robinson, S.R., E.C.G.M. Hampson, M.N. Munro, and D.I. Vaney. 1993. Unidirectional coupling of gap junctions between neuroglia. Science 262:1072–1074.
Stems are from the chapter Problems section; correct choices are drawn from Abbreviated Solutions to Problems (Appendix B) in the same edition.
1. Entropy Changes during Egg Development Consider a system consisting of an egg in an incubator. The white and yolk of the egg contain proteins, carbohydrates, and lipids. If fertilized, the egg transforms from a single cell to a complex organism. Discuss this irreversible process in terms of the entropy changes in the system and surroundings. Be sure that you first clearly define the system and surroundings.
2. Calculation of ΔG′° from an Equilibrium Constant Calculate the standard free-energy change for each of the three metabolically important enzyme-catalyzed reactions, using the equilibrium constants given for the reactions at 25 °C and pH 7.0. a. b. c.
3. Calculation of the Equilibrium Constant from ΔG′° Calculate the equilibrium constant K′eq for each of the three reactions at pH 7.0 and 25 °C, using the ΔG′° values in Table 13-4. a. b. c.
4. Experimental Determination of K′eq and ΔG′° Incubating a 0.1 M solution of glucose 1-phosphate at 25 °C with a catalytic amount of phosphoglucomutase transforms some of the glucose 1-phosphate to glucose 6-phosphate. At equilibrium, the concentrations of the reaction components are G lucose 1-phosphate ⇌ glucose 6-phosphate 4.5× 10−3M 9.6× 10−2M Calculate K′eq and ΔG′° for this reaction.
5. Experimental Determination of ΔG′° for ATP Hydrolysis A direct measurement of the standard free-energy change associated with the hydrolysis of ATP is technically demanding because the minute amount of ATP remaining at equilibrium is difficult to measure accurately. The value of ΔG′° can be calculated indirectly, however, from the equilibrium constants of two other enzymatic reactions having less favorable equilibrium constants: G lucose 6-phosphate+ H2O → glucose+ Pi K′eq = 270 AT P + glucose→ AD P + glucose 6-phosphate K′eq = 890 Using this information for equilibrium constants determined at 25 °C, calculate the standard free energy of hydrolysis of ATP.
6. Difference between ΔG′° and Δ G Consider the interconversion shown, which occurs in glycolysis (Chapter 14): Fructose 6-phosphate ⇌ glucose 6-phosphate K′eq = 1.97 a. What is ΔG′° for the reaction (K′eq measured at 25 °C)? b. If the concentration of fructose 6-phosphate is adjusted to 1.5 M and that of glucose 6-phosphate is adjusted to 0.50 M, what is ΔG? c. Why are ΔG′° and ΔG different?
7. Free Energy of Hydrolysis of CTP Compare the structure of the nucleoside triphosphate CTP with the structure of ATP. Now predict the K′eq and ΔG′° for the reaction: AT P + CD P → AD P + CT P
8. Dependence of Δ G on pH The free energy released by the hydrolysis of ATP under standard conditions is −30.5 kJ/mol. If ATP is hydrolyzed under standard conditions except at pH 5.0, is more or less free energy released? Explain.
9. The ΔG′° for Coupled Reactions Glucose 1-phosphate is converted into fructose 6-phosphate in two successive reactions: G lucose 1-phosphate→ glucose 6-phosphate G lucose 6-phosphate→ fructose 6-phosphate Using the ΔG′° values in Table 13-4, calculate the equilibrium constant, K′eq, for the sum of the two reactions: G lucose 1-phosphate → fructose 6-phosphate
10. Effect of [ATP]/[ADP] Ratio on Free Energy of Hydrolysis of ATP Using Equation 13-4, plot ΔG against ln Q (mass-action ratio) at 25 °C for the concentrations of ATP, ADP, and Pi in the table shown. ΔG′° for the reaction is −30.5 kJ/mol. Use the resulting plot to explain why metabolism is regulated to keep the ratio [ATP]/[ADP] high. Concentration (m ) ATP 5 3 1 0.2 5 ADP 0.2 2.2 4.2 5.0 25 Pi 10 12.1 14.1 14.9 10
11. Strategy for Overcoming an Unfavorable Reaction: ATP- Dependent Chemical Coupling The phosphorylation of glucose to glucose 6-phosphate is the initial step in the catabolism of glucose. The direct phosphorylation of glucose by Pi is described by the equation a. Calculate the equilibrium constant for this reaction at 37 °C. In the rat hepatocyte, the physiological concentrations of glucose and Pi are maintained at approximately 4.8 mM. What is the equilibrium concentration of glucose 6-phosphate obtained by the direct phosphorylation of glucose by Pi? Does this reaction represent a reasonable metabolic step for the catabolism of glucose? Explain. b. In principle at least, one way to increase the concentration of glucose 6-phosphate is to drive the equilibrium reaction to the right by increasing the intracellular concentrations of glucose and Pi. Assuming a fixed concentration of Pi at 4.8 mM, how high would the intracellular concentration of glucose have to be to give an equilibrium concentration of glucose 6-phosphate of 250 μ M (the normal physiological concentration)? Would this route be physiologically reasonable, given that the maximum solubility of glucose is less than 1 M? c. The phosphorylation of glucose in the cell is coupled to the hydrolysis of ATP; that is, part of the free energy of ATP hydrolysis is used to phosphorylate glucose: G lucose+ Pi → glucose 6-phosphate+ H2O ΔG′°= 13.8 kJ/m Calculate K′eq at 37 °C for the overall reaction. For the ATP-dependent phosphorylation of glucose, what concentration of glucose is needed to achieve a 250 μ M intracellular concentration of glucose 6-phosphate when the concentrations of ATP and ADP are 3.38 mM and 1.32 mM, respectively? Does this coupling process provide a feasible route, at least in principle, for the phosphorylation of glucose in the cell? Explain. d. Although coupling ATP hydrolysis to glucose phosphorylation makes thermodynamic sense, we have not yet specified how this coupling is to take place. Given that coupling requires a common intermediate, one conceivable route is to use ATP hydrolysis to raise the intracellular concentration of Pi and thus drive the unfavorable phosphorylation of glucose by Pi. Is this a reasonable route? (Think about the solubility product, Ksp, of metabolic intermediates.) e. The ATP-coupled phosphorylation of glucose is catalyzed in hepatocytes by the enzyme glucokinase. This enzyme binds ATP and glucose to form a glucose-ATP-enzyme complex, and the phosphoryl group is transferred directly from ATP to glucose. Explain the advantages of this route. (1) G lucose+ Pi→ glucose 6-phosphate+ H2O ΔG′°= 13.8 kJ/mo (2) AT P + H2O → AD P + Pi ΔG′°= −30.5 kJ––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– Sum : G lucose+ AT P → glucose 6-phosphate+ AD
12. Calculations of ΔG′° for ATP-Coupled Reactions From data in Table 13-6, calculate the ΔG′° value for each reaction: a. Phosphocreatine+ AD P → creatine+ AT P b. AT P + fructose→ AD P + fructose 6-phosphate
13. Coupling ATP Cleavage to an Unfavorable Reaction To explore the consequences of coupling ATP hydrolysis under physiological conditions to a thermodynamically unfavorable biochemical reaction, consider the hypothetical transformation X → Y, for which ΔG′°= 20.0 kJ/mol. a. What is the ratio [Y]/[X] at equilibrium? b. Suppose X and Y participate in a sequence of reactions during which ATP is hydrolyzed to ADP and Pi. The overall reaction is X + AT P + H2O → Y + AD P + Pi Calculate [Y]/[X] for this reaction at equilibrium. Assume that the temperature is 25.0 °C and the equilibrium concentrations of ATP, ADP, and Pi are 1 M. c. We know that [ATP], [ADP], and [Pi] are not 1 M under physiological conditions. Calculate [Y]/[X] for the ATP- coupled reaction when the values of [ATP], [ADP], and [Pi] are those found in rat myocytes (Table 13-5).
14. Calculations of Δ G at Physiological Concentrations Calculate the actual, physiological ΔG for the reaction Phosphocreatine+ AD P → creatine+ AT P at 37 °C, as it occurs in the cytosol of neurons, with phosphocreatine at 4.7 mM, creatine at 1.0 mM, ADP at 0.73 mM, and ATP at 2.6 mM.
15. Free Energy Required for ATP Synthesis under Physiological Conditions In the cytosol of rat hepatocytes, the temperature is 37 °C and the mass-action ratio, Q, is = 5.33× 102 M −1 Calculate the free energy required to synthesize ATP in a rat hepatocyte.
16. Chemical Logic In the glycolytic pathway, a six-carbon sugar (fructose 1,6-bisphosphate) is cleaved to form two three-carbon sugars, which undergo further metabolism. In this pathway, an isomerization of glucose 6-phosphate to fructose 6-phosphate (as shown in the diagram) occurs two steps before the cleavage reaction. The intervening step is phosphorylation of fructose 6- phosphate to fructose 1,6-bisphosphate (p. 516). [AT P] [AD P][Pi] What does the isomerization step accomplish from a chemical perspective? (Hint: Consider what might happen if the C—C bond cleavage were to proceed without the preceding isomerization.)
17. Enzymatic Reaction Mechanisms I Lactate dehydrogenase is one of the many enzymes that require NADH as coenzyme. It catalyzes the conversion of pyruvate to lactate: Draw the mechanism of this reaction (show electron-pushing arrows). (Hint: This is a common reaction throughout metabolism; the mechanism is similar to that catalyzed by other dehydrogenases that use NADH, such as alcohol dehydrogenase.)
18. Enzymatic Reaction Mechanisms II Biochemical reactions o en look more complex than they really are. In the pentose phosphate pathway (Chapter 14), sedoheptulose 7-phosphate and glyceraldehyde 3-phosphate react to form erythrose 4- phosphate and fructose 6-phosphate in a reaction catalyzed by transaldolase. Draw a mechanism for this reaction (show electron-pushing arrows). (Hint: Take another look at aldol condensations, then consider the name of this enzyme.)
19. Recognizing Reaction Types For each pair of biomolecules, identify the type of reaction (oxidation-reduction, hydrolysis, isomerization, group transfer, or internal rearrangement) required to convert the first molecule to the second. In each case, indicate the general type of enzyme and cofactor(s) or reactants that would be required, and any other products that would result. a. b. c. d. e. f. g.
20. Effect of Structure on Group Transfer Potential Some invertebrates contain phosphoarginine. Is the standard free energy of hydrolysis of this molecule more similar to that of glucose 6-phosphate or of ATP? Explain your answer.
21. Polyphosphate as a Possible Energy Source The standard free energy of hydrolysis of inorganic polyphosphate (polyP) is about −20 kJ/mol for each Pi released. We calculated in Worked Example 13-2 that, in a cell, it takes about 50 kJ/mol of energy to synthesize ATP from ADP and Pi. Is it feasible for a cell to use polyphosphate to synthesize ATP from ADP? Explain your answer.
22. Daily ATP Utilization by Human Adults a. The synthesis of ATP from ADP and Pi requires a total of 30.5 kJ/mol of free energy when the reactants and products are at 1 M concentrations and the temperature is 25 °C (standard state). However, the actual physiological concentrations of ATP, ADP, and Pi are not 1 M, and the physiological temperature is 37 °C. Thus, the free energy required to synthesize ATP under physiological conditions is different from ΔG′°. Calculate the free energy required to synthesize ATP in the human hepatocyte when the physiological concentrations of ATP, ADP, and Pi are 3.5, 1.50, and 5.0 mM, respectively. b. A 68 kg (150 lb) adult requires a caloric intake of 2,000 kcal (8,360 kJ) of food per day (24 hours). The body metabolizes the food and uses the free energy to synthesize ATP, which then provides energy for the body’s daily chemical and mechanical work. Assuming that the efficiency of converting food energy into ATP is 50%, calculate the weight of ATP used by a human adult in 24 hours. What percentage of the body weight does this represent? c. Although adults synthesize large amounts of ATP daily, their body weight, structure, and composition do not change significantly during this period. Explain this apparent contradiction.
23. Rates of Turnover of γ and β Phosphates of ATP A er adding a small amount of ATP labeled with radioactive phosphorus in the terminal position, [γ-32P]ATP, to a yeast extract, a researcher finds about half of the 32P activity in Pi within a few minutes, but the concentration of ATP remains unchanged. Explain. She then carries out the same experiment using ATP labeled with 32P in the central position, [β-32P]ATP, but the 32P does not appear in Pi within such a short time. Why?
24. Cleavage of ATP to AMP and PPi during Metabolism Synthesis of the activated form of acetate (acetyl-CoA) is carried out in an ATP-dependent process: Acetate+ CoA + AT P → acetyl-CoA + AM P + PPi a. The ΔG′° for hydrolysis of acetyl-CoA to acetate and CoA is −32.2 kJ/mol. The ΔG′° for hydrolysis of ATP to AMP and PPi is −30.5 kJ/mol. Calculate ΔG′° for the ATP- dependent synthesis of acetyl-CoA. b. Almost all cells contain the enzyme inorganic pyrophosphatase, which catalyzes the hydrolysis of PPi to Pi. What effect does the presence of this enzyme have on the synthesis of acetyl-CoA? Explain.
25. Activation of a Fatty Acid by Reaction with Coenzyme A In the reaction sequence for fatty acid breakdown, coenzyme A (CoA), with its thiol (— SH) group, joins to the fatty acid as a thiol ester, as ATP is converted into AMP and PPi: R—COO− + AT P + CoA—SH → AM P + PPi+ R—CO—S—CoA The oxidation of fatty acids as fuels requires two steps. The first step transfers an activating group from ATP to the carboxyl group of the fatty acid. In the second step, CoA—SH displaces the activating group to form fatty acyl-S—CoA. Given the known products of the reaction, what is the activating group?