BOTANY VOLUME 2 - PLANT PHYSIOLOGY - 2007

6. PHYSIOLOGY OF METABOLISM

Life processes are intimately linked to the continuous transformation of matter and energy. Living organisms take up specific substances and energy from their environment and release other substances and energy (particularly heat) back into it. In Thermodynamics (Greek therme, heat; dynamis, driving force), such systems are referred to as open systems. Ultimately, the vast majority of the energy introduced into the biosphere originates from sunlight, which green plants convert into chemical energy through Photosynthesis. This process transforms inorganic substances into Organic compounds. Organisms that synthesize all necessary organic compounds from inorganic precursors are called autotrophs (primary producers). If plants utilize light energy, they are termed photoautotrophs. Some microorganisms live chemoautotrophically, meaning they utilize both the matter and the energy of Inorganic Compounds. Heterotrophic organisms (consumers) derive their matter and energy from primary producers. Consequently, they depend on the organic compounds synthesized by primary producers and satisfy their Energy Requirements by consuming organic matter. Among heterotrophs, Saprophytes¹ feed on non-living nutrient sources, whereas parasites obtain nourishment from living organisms (Table 6.1; see 9.1.1).

¹ Among vascular plants, a purely saprophytic nutritional strategy does not occur. As a rule, when utilizing non-living organic sources, vascular plants form symbioses or parasitize Fungi. — Translator's Note.

The transformation of matter and energy within The Cell, known as METABOLISM (Greek metabole, change, alteration), can be divided into anabolic (biosynthetic) and catabolic (degradative) processes. The Metabolic pathways of fundamental importance to life Functions constitute primary metabolism. Plants are particularly notable for their richly differentiated Secondary Metabolism. Secondary metabolism comprises SPECIALIZED METABOLIC PATHWAYS that branch off from primary metabolites (hence designated as secondary based solely on this origin rather than any lesser significance) and lead to products with auxiliary functions—often ecological, such as defense compounds against herbivores. Secondary metabolites are largely restricted to specific plant groups and thus hold taxonomic value.

Class="center">Table 6.1. Various pathways of carbon assimilation in organisms

The first section of this chapter will initially examine the fundamental thermodynamic principles of life processes (see

6.1), followed by the autotrophic functions of plants, beginning with the uptake and Processing of mineral nutrients (6.2), which are closely linked to Water relations (6.3). The synthesis of organic compounds from absorbed inorganic precursors using light energy (photosynthesis) and the distribution of photosynthetic products (assimilates) within the plant comprise two sections (6.4, 6.5) covering plant primary metabolism (6.4–6.15), which is followed by An Overview of key aspects of secondary metabolism (6.16) and the metabolism of plant-specific polymers (6.17). The chapter concludes with a brief description of excretory processes in plants (6.18).

6.1. Energetics of Metabolism

6.1.1. Fundamentals of Bioenergetics

There is no doubt that the transformations of matter and energy in living organisms obey the laws of physics and chemistry, and that the principles of thermodynamics—the science of Energy Changes During physical or chemical processes—apply equally to living beings. While energy transformations in living Cells are often collectively termed bioenergetics (Greek energia, action), this simply signifies that among thermodynamically feasible processes and transformations, certain ones are particularly characteristic of the living cell, and that the types of molecules involved in these reactions, especially catalysts, differ from those found in inanimate nature and technology.

Living organisms are open systems (in the thermodynamic sense), meaning they are in a constant exchange of energy and matter with their environment. They evolve, which implies that their material and energy exchange is subject to temporal changes. Furthermore, life processes are irreversible, and a living Organism is far removed from a state of thermodynamic equilibrium. Consequently, living beings should be described using the principles of irreversible nonequilibrium thermodynamics; however, given the staggering complexity of biological processes, this currently appears to be an unrealistic undertaking¹. Significant fundamental insights can already be gained from the much simpler equilibrium thermodynamics of closed systems—systems that exchange energy, but not matter, with their environment. This information helps determine whether a specific chemical reaction is feasible under given conditions. Nevertheless, the laws of equilibrium thermodynamics say nothing about the rate at which a reaction will proceed.

¹ A complete description of all thermodynamic processes within a cell is fundamentally impossible because the cell is coupled with its external environment. It is always possible to introduce a new environmental factor that alters cell metabolism, rendering the description incomplete. — Translator's note.

The metabolism of a living cell serves to perform specific functions and carry out work that requires energy.

The absolute measure of work (force × distance), much like energy, is the Joule (J; 1 kg 1 m² 1 s^-2 = kg m² s^-2), which is the unit of force (Newton, N: kg m s^-2) × unit of distance (m). Data are frequently expressed in kilojoules (kJ; 10³ J). A previously used and still common measure of energy is the calorie (1 cal = 4.184 J), a measure of heat (1 cal corresponds to The amount of energy required to raise the Temperature of 1 g of water at standard pressure from 14.5 to 15.5 °C). The validity of using this unit as a universal measure of energy stems from the mutual interconvertibility of different forms of energy, i.e., the transition of one form into another (such as kinetic, thermal, chemical, electrical, and radiant energy). The preference for the heat unit was based on the fact that heat is the most general form of energy; all Other forms of energy can be completely converted into heat (but not vice versa). Temperature is most commonly expressed in degrees Celsius (°C), although it is more correct to use absolute temperature in kelvins, K (0 K = -273.15 °C). (A table of SI units and conversion factors is provided at the end of the book.)

6.1.2. Energetics of Closed Systems

Thermodynamics typically examines the behavior (more precisely, the state change, A) of a delimited region (system). Everything outside the system constitutes its environment (surroundings). The System and Its environment together are referred to as the "whole," the "universe," or "universum" (Fig. 6.1). A system possesses internal energy U, which represents the sum of all forms of energy within the system. The first law of thermodynamics states that the internal energy of a closed system—that is, a system that exchanges neither matter nor energy with its environment—remains constant (U = const). The amount of energy in a system depends on its state, not on the path by which that state was reached. Therefore, for a cyclic process in which the system returns to its initial state, ΔU = 0. Thus, energy can neither be created nor destroyed.

Fig. 6.1. Structure/97.html">Definitions of various thermodynamic systems

If energy is introduced into the system from the outside, such as a specific amount of heat (Q) (making it, by definition, not an isolated system but a closed system, since it exchanges energy, but not matter, with the surroundings), then, According to the first law, the introduced heat leads either to A change in the internal energy of the system or to the performance of work (W):

Q = ΔU + W or ΔU = Q - W. (6.1)

Processes in which a system absorbs heat are called endothermic, whereas processes that release heat are termed exothermic. In reactions at constant pressure (p = const), as typically occurs in living organisms, the heat change is referred to as the Enthalpy Change and is denoted by ΔH (Q = ΔH). In this case,

ΔU = ΔH - W, (6.2)

where W generally takes the form of volume-change work: W = pΔV. At constant volume and constant pressure, therefore, no work is performed (W = 0), and

∆U = ∆Н.

Under these conditions, measuring the heat of reaction allows one to draw Conclusions about the energy change accompanying the process. The enthalpy change (∆H) of a reaction can be determined by calorimetry (from Greek calor, heat, and metrein, to measure). If ∆H > 0, the process is referred to as endothermic, whereas if ∆H > 0, it is called exothermic. Organic compounds possess a specific molar Heat of Combustion, defined as the energy (in joules) released into the environment upon the Complete oxidation of 1 mole of a substance (Table 6.2).

Table 6.2. Heat of combustion of various organic compounds important for cellular metabolism

The First Law of thermodynamics does not allow us to predict the direction of physical or chemical processes. However, general experience tells us that spontaneously (i.e., autonomously) occurring processes have a definite direction. For instance, heat flows from a warmer body to a cooler one; the reverse process has never been observed. In general, only states of lower Organization arise spontaneously from states of higher organization, considering the system and its surroundings as a whole. The measure of disorder is the thermodynamic function S, known as Entropy (from Greek entrepein, to turn about). Every spontaneous change in state is accompanied by an increase in entropy. This is one of the formulations of the second law of thermodynamics. A protein molecule that spontaneously transitions from an unfolded conformation with a low degree of organization to a folded, highly organized state during the Formation of secondary and tertiary structures might seem to contradict this principle. However, the folding process is accompanied by the disruption of the water structure surrounding the folding protein molecule, so that the total entropy of the system (the protein) and its environment (the aqueous medium) increases during folding. Similarly, maintaining a state of high organization (low entropy) in living organisms is inextricably linked to an increase in entropy in their environment.

The dimension of entropy is [J K^-1]. At any given temperature, solids have a relatively low entropy, liquids an intermediate one, and gases a high entropy. Entropy increases with temperature because thermal molecular motion increases. The entropy of a perfectly crystalline substance at absolute zero (-273.15 °C = 0 K) is zero (a principle often referred to as the third law of thermodynamics).

As already mentioned, an influx of heat into a system can be used to perform work, as occurs, for example, in heat engines. In a living cell, however, temperature remains virtually constant¹: the cell functions under nearly isothermal conditions. The fraction of the total enthalpy of a system capable of performing work under isothermal conditions is designated as free enthalpy (G) (commonly known as Gibbs Free energy). The fundamental equation relating the changes in entropy and enthalpy to The change in free enthalpy is as follows:

∆G = ∆Н - Т ∆S. (6.3)

¹ In some cases, temperature fluctuations are significant, such as during thermogenesis in the inflorescences of arum lilies. The temperature increase is associated with exergonic Respiration reactions (see below). — Editor's note.

Here, ∆G represents the change in the free enthalpy of the system; the enthalpy change ∆Н is the heat exchanged between the system and its surroundings when the system performs no work (see above); T is the absolute temperature (in K); and ∆S is the entropy change of the system.

The sign of ∆G determines whether a given reaction can proceed spontaneously. If ∆G > 0, the reaction is not spontaneous and can only occur if ∆G < 0 (though not necessarily rapidly). A spontaneously occurring reaction proceeds with a decrease in free enthalpy and an increase in entropy until ∆Н balances T ∆S, thereby reaching a state where ∆G = 0 (equilibrium state). Processes with ∆G < 0 are termed exergonic, whereas those with ∆G > 0 are called endergonic.

At T = 0 or ∆S = 0, the ∆G of a reaction can be determined from the thermal effect of the reaction (i.e., the enthalpy change ∆Н). However, these conditions are not met in biological systems. Nevertheless, ∆G and consequently the driving force of a reaction can be approximately determined from its thermal effect when the magnitude of ∆Н is large—such as during The oxidation of nutrients in respiration (see 6.10.3)—and the temperature is low (as is often the case in cells), so that the T∆S term has little influence on the free enthalpy value. Conversely, in processes with a small thermal effect, such as hydrolytic Cleavage and polymerization (Condensation)—which are equally vital biological reactions—the entropy change can significantly determine the free enthalpy. Because ∆S is coupled with Temperature, The Influence of entropy on ∆G increases in proportion to the absolute temperature.

To understand the course of Chemical Reactions, it is useful to relate the free enthalpy of a reaction to changes in amount of substance and, ultimately, to the resulting chemical equilibrium. This is supported by the fact that a chemical reaction A → B, accompanied by a decrease in free energy, will proceed until the enthalpy minimum (∆G = 0) is reached. Following this, no further net change in the amounts of substances occurs, an equilibrium A ⇄ B is established, and The ratio of product concentrations to reactant concentrations remains constant (law of mass action). This ratio is called the thermodynamic Equilibrium Constant, K:

(6.4)

and in the general case for the reaction A + B ⇄ C + D:

(6.5)

Thus, the equilibrium constant is expressed as the product of the concentrations of the reaction products divided by the product of the concentrations of the reactants at equilibrium. The relationship between ∆G and K is expressed by the equation

∆G⁰ = RТ ln К (J • mol^-1), (6.6)

where ∆G⁰ is the change in standard molar free enthalpy (per mole of converted substance under standard conditions: temperature T = 25 °C and p = 1 atm = 0.1 MPa), T is the absolute temperature in Kelvin (K), and R is the universal gas constant (= 8.314 J • mol^-1 • K^-1).

For reactions involving hydrogen ions—as is frequently the case in biological systems—their standard concentration must also be set to 1 mole. For practical reasons, biochemical literature defines standard conditions by specifying not the amount of substance in moles, but rather the change in molar concentration (mol L^-1). This means that under standard conditions, the hydrogen ion concentration change would be 1 mol • L^-1 (pH 0), which is a completely non-physiological value. Therefore, biochemical literature uses a slightly modified Definition of the standard state, in which the H⁺ ion concentration is set to 10^-7 M (pH 7), and the concentration of water (55.5 mol • L^-1), which remains practically constant during the reaction, is incorporated into the constants (if water appears in the reaction equation):

∆G⁰' = RТ ln К'.   (6.7)

Table 6.3. Changes in standard molar free enthalpy at pH 7 (∆G⁰') for some important metabolic reactions (Hydrolysis reactions)

¹ In some literature, inorganic phosphate is denoted as P_i (from inorganic). — Editor's note.

Changes in standard molar free enthalpy (at pH 7) for several important reactions are listed in Table 6.3.

Inside the cell, however, conditions prevail that differ significantly from standard ones. For instance, the pH often deviates from 7.0, the temperature from 25 °С, and substance concentrations, as a rule, do not match standard conditions. Therefore, a careful distinction must be made between the change in standard molar Gibbs free energy ∆G0', which remains constant at a given temperature, and the actual change in free energy ∆G, which depends on the actual temperature and the real concentrations of the reaction components. It is ∆G, rather than ∆G⁰', that determines the direction of a reaction within the cell. In many cases, however, determining these ∆G values is extremely difficult because the actual conditions (substance concentrations, pH values, temperature) within individual reaction spaces (compartments) are hard to measure precisely.

6.1.3. Energetics of Open Systems

While the thermodynamics of closed or equilibrium systems yields important conclusions regarding the energetics of individual biochemical reactions (such as whether a given process can occur spontaneously or not), living organisms are open systems that continuously exchange energy and matter with their environment (see Fig. 6.1). Whereas any closed system tends toward a stationary equilibrium state (∆G = 0), open systems are able to maintain a stable state far removed from thermodynamic equilibrium, namely a steady state. The thermodynamic description of such open systems falls under non-equilibrium or irreversible thermodynamics, which primarily accounts for the Time Factor and in which matter fluxes play a major role. Although a detailed Discussion of irreversible thermodynamics is beyond the scope here, THE CONCEPT OF chemical potential (see 6.1.4) proves extremely useful for better understanding the energetics of many physiological processes.

A stable steady state is characterized by the continuous flow of matter and energy through the system, which constantly generates free energy within it. Ultimately, this occurs via the exergonic conversion of high-energy, low-entropy organic compounds (nutrients) into low-energy, high-entropy "waste products" (see Fig. 6.10). Photosynthetically active cells produce these products (primary production) primarily from inorganic compounds using absorbed light energy in a highly exergonic photosynthesis process (see 6.4–6.7). Free energy in the form of energy-rich compounds, such as ATP, is utilized to perform biological work and maintain the high level of organization characteristic of living beings. If the flow of matter and energy is interrupted, a state of equilibrium (∆G = 0) is established after some time—resulting in death.

It has been demonstrated that a steady state is a state of an open system in which entropy production is minimized and the maximum possible level of organization is maintained at a minimum energy cost. Thus, the steady state represents the condition of an open system with maximum thermodynamic efficiency. Crucially, unlike the static equilibrium established in a closed system, a system in a steady state can be regulated (an essential property of all living cells).

6.1.4. Chemical Potential

6.1.4.1. General Definition

The Free energy of an open system with a complex composition, such as a living cell, is practically impossible to determine. However, in many cases, it is sufficient to establish the capacity of specific components within this system to perform work. For example, of interest is only the free energy difference of protons (rather than other ions) across The cell membrane, provided one can calculate the proton motive force driving work in coupled transport processes and determine its direction. Alternatively, It is interesting to evaluate the difference in the free energy of water in adjacent aqueous solutions separated by a cell membrane, as well as the direction and scale of water flow across this interface.

The free energy calculated per mole of component i in a mixture of k components is called the chemical potential μ of component i (μ_i). The sum of the chemical potentials of all k components yields the free energy per mole of the substance mixture. Thus, the contributions of the individual components Complement one another. The chemical potential of each component in a substance mixture can, in turn, be broken down into a standard potential (μ_i⁰) and a sum of terms reflecting deviations from the standard state:

where RT ln x₁ is the concentration term: R is the universal gas constant; T is the absolute temperature; x₁ is the mole fraction of i (x₁ = n₁ : (n_a + n_b ... n_k)). The mole fraction is the ratio of the amount of substance (in moles) of a given component to the total amount of all substances present in the solution, including the solvent; pV₁ is the pressure term: p is pressure, V₁ is the partial molar volume of i, corresponding to the change in the system's volume upon The addition of 1 mole of component i; ghM₁ is the gravitational term: g is the gravitational constant (9.806 m s^-2), h is the elevation height, M₁ is the molar mass of i; FEz₁ is the electrical term: F is the Faraday constant (96.49 kJ • V^-1 • mol^-1), E is the electrical potential, z₁ is the valence charge of i.

The dimensions of μ are expressed in units of energy per mole (J • mol^-1).

Since one is often interested not in the chemical potential itself, but in its change when the state of the system with respect to component i changes, the relation for calculating the change in chemical potential (= free energy) of i in a mixture during a state transition A → B is derived as follows:

Variants of general equations 6.8 and 6.9 are relevant for subsequent chapters and will be discussed below.

6.1.4.2. Water Potential

Since plant cells, like those of other organisms, cannot actively transport water, it moves passively from a region of higher (more positive) free energy to a region of lower (more negative) free energy—meaning this is an exergonic process that occurs spontaneously (though not necessarily rapidly). Because biological systems essentially involve aqueous mixtures with other substances (e.g., aqueous solutions in cells and soil, water vapor in the gas phase of the atmosphere), it is practical to use the concept of water potential (μ_H2O) in energy calculations. Water molecules are electrically uncharged (z_H2O = 0), so the electrical term drops out of equation 6.9, reducing it to

Consequently, for pure water (x_H2O = 1) in the standard state (p = 0, h = 0), the value is μ_H2O = μ⁰_H2O.

Based on the relation x_H2O = 1 - Σ_ix_i, the concentration term RT ln x_H2O can be expressed as a function of the mole fraction of all dissolved solutes: RT ln (1 - Σ_ix_i). For dilute solutions, the approximation ln (1 - x) = -x can be applied; using the relation Σ_ix_i = V_H2O Σ_ic_i, where c is the molar concentration, ultimately yields the equation

For more concentrated solutions (generally 0.1 M and above), molal concentrations (mol • kg^-1) and activities should be used instead of concentrations.

Since RT Σ_ic_i = Π (Π being the osmotic pressure, van 't Hoff's law), this gives

and hence

Since in this case as well, the difference in chemical potential of water is generally of greater interest than the absolute value of the potential, subsequent transformations regarding the partial molar volume of water primarily determine the deviation of the water's chemical potential in the system under consideration from its standard state

as the water potential of the solution. Furthermore, equation 6.12 implies that

has the dimensions of energy/volume (=force/area = pressure) and is expressed in units of bar or Pa (1 bar = 0.1 MPa).

On a cellular scale, height differences are negligible, so equation 6.14 simplifies further upon omitting the gravitational component to

The water potential of a solution, i.e., the free enthalpy of water relative to the partial molar volume of water (V_Н2O - 18 mL), is therefore determined by three component potentials:

✵ pressure potential, p (the hydrostatic pressure at which the solution is stationary);

✵ osmotic potential, -П (the negative value of the osmotic pressure П);

✵ gravitational potential (the latter can be neglected when considering processes on a cellular scale).

It should be noted that hydrostatic pressure is defined as the deviation from ambient pressure. It can take both positive values ("excess pressure") and negative values ("tension", "suction"). Absolute pressure is always positive and equals zero in a complete vacuum, respectively. Consequently, the water pressure potential (p) in the standard state is equal to 0 (p = 0), and its absolute pressure is 1 bar (0.1 MPa).

If There is a difference between the water potentials of two compartments (∆'=0), water will continuously flow from the region with a more positive water potential to the region with a more negative water potential. This process decreases its free enthalpy, making it exergonic and therefore spontaneous.

The concept of water potential and its implications prove extremely useful for understanding plant water relations as a whole (see 6.3).

6.1.4.3. Chemical Potential of Ions and Transmembrane Potential

The chemical potential of electrically charged particles in a solution is determined primarily by their concentration and electrical charge. Accordingly, the chemical potential equation for ion i is written as:

(a₁ is The activity of ion i; for dilute solutions a₁∞ c₁; c₁ is the molar concentration of i).

If we consider two solutions of i in compartments A and B separated by an electrically insulating membrane, the difference in chemical potential of i, ∆μ1 (also referred to as the Electrochemical Potential), is defined as

The electrochemical potential of hydrogen ions across cell membranes will be of particular importance hereafter, as it serves as the driving force for many transport processes across cell membranes, for ATP Synthesis in Chloroplasts (see 6.4.9), and in Mitochondria (see 6.10.3.3). For H⁺, Z_H+ = 1, and hence we obtain

Here, reaction space A represents the intracellular compartment, and reaction space B represents the extracellular (or functionally extracellular) compartment. The potential difference E^B - E^A = ∆Е_М is designated as the transmembrane electrical potential (briefly: Membrane Potential). In a simplified form, combining all constants at a standard temperature (T = 298 K) and applying the definition of pH (pH = -log[Н⁺]), we obtain

¹ Here and below, log denotes the common logarithm. In German literature, it is usually denoted by the symbol lg. — Translator's note.

The expression is referred to as the proton motive force (pmf) and is used to characterize the energy of a proton gradient. The two components, either individually or jointly, are capable of performing work: on the one hand, the hydrogen ion concentration potential (∆pH), and on the other hand, the electrical potential (∆E_M). For Examples, see 6.1.5.

For The equilibrium state (∆μ₁ = 0), equation 6.17 yields

This equation is called the Nernst equation (∆E_N, Nernst equilibrium potential, V).

For , a potential difference arises between the two compartments. With an effective 10-fold concentration difference (at z₁ = 1), the voltage difference will be 59 mV (at z₁ = 2, respectively, 29.5 mV). Conversely, with an applied constant voltage of 59 mV at equilibrium, a concentration difference of 1:10 will be established between these compartments for a permeable ion.

6.1.4.4. Redox Potential

Numerous biologically relevant metabolic transformations involve the reduction or oxidation of metabolites. Reduction is defined as the gain of electrons, whereas oxidation refers to the loss of electrons by a molecule. Oxidation and reduction generally proceed as coupled processes (oxidation-reduction or redox reactions). Redox reactions can also be described using the chemical potential (electrochemical potential) defined for ions in equation 6.17. Thus, for the coupled reactions A_ox + B_red ⇄ A_red + B_ox, the Nernst equation takes the following form:

where R, T, and F are the previously introduced quantities; z is the number of electrons transferred according to the reaction equation; ∆E⁰ is the difference between the standard redox potentials of the oxidizing and reducing agents: ∆E⁰ = E⁰_B - E⁰_A. These are determined for the reductant and oxidant as the potential difference relative to the standard hydrogen electrode under standard conditions (whose potential is assumed to be 0) and thus themselves represent a potential difference.

E⁰ or ∆E⁰ values are conventionally standardized to 25 °C, a pressure of 1 atm (0.1 MPa), and a substance concentration change of 1 mol • L^-1. If hydrogen ions (protons) participate in redox reactions, there is a corresponding change in amount of substance of 1 mol/L (pH 0, see equation 6.6). For biological purposes, owing to these circumstances, a different definition of standard conditions (E⁰) is chosen here as well, analogously to ∆G⁰ earlier: pH 7. There exists the relationship

E⁰' = E⁰ - 0.42 V. (6.22)

Some standard potentials for pH 7 are summarized in Table 6.18 (see 6.4.5).

The redox potential ∆E expresses the electrochemical energy supplied by a redox reaction to perform work per transferred mole of electrons. The Gibbs free energy change of a reaction can be readily determined from the redox potential via the relation:

∆G = -zF • ∆E.    (6.23)

Accordingly,

∆G⁰' = -zF • ∆E⁰'.   (6.24)

Standard redox potentials can be used to determine the direction in which coupled redox reactions will proceed spontaneously (though not necessarily rapidly) under standard conditions. A redox reaction is exergonic (∆G < 0) if electrons are transferred from a reaction partner with a more negative standard redox potential (i.e., the reductant, which is oxidized in the course of the reaction) to a reaction partner with a more positive standard redox potential (i.e., the oxidant, which is reduced during the reaction). However, since standard conditions do not prevail within the cell, analyzing ∆E⁰ values does not necessarily reflect the actual course of a reaction in vivo. This requires knowledge of ∆E (and hence ∆G), which fundamentally entails knowing the actual concentrations of the components involved in the redox reaction, as well as the actual temperature and pH. These factors, however, are generally not known with absolute precision; therefore, standard values are frequently used to evaluate the fundamental energetic relationships of redox reactions, as is standard practice in biochemical processes in general.

Redox reactions play a central role in metabolism. Both photosynthesis and cellular respiration are redox processes (Fig. 6.2). During photosynthesis, carbon is reduced from the oxidation state of CO₂ (oxidation number +4) to the reduction level of CARBOHYDRATES ([CH₂O]_n, oxidation number 0). Electrons are derived from water and, via a complex light-driven (endergonic) Electron Transport Chain, are initially transferred to oxidized nicotinamide adenine dinucleotide phosphate (NADP⁺) to form NADPH, which serves as a carrier molecule for reducing equivalents and is subsequently re-oxidized in CO₂ assimilation reactions (see 6.5.2). Mitochondrial respiration also involves the oxidation of carbohydrates to CO₂ in order to channel the resulting electrons into another widespread carrier molecule of redox equivalents—reduced nicotinamide adenine dinucleotide (NADH), derived from oxidized nicotinamide adenine dinucleotide (NAD⁺)—followed by electron transport via membrane complexes to oxygen (see 6.10.3.3). In addition to these two fundamental metabolic redox processes, numerous other oxidations and reductions of metabolites catalyzed by redox Enzymes (oxidoreductases) play crucial roles in metabolism.

Fig. 6.2. Energetic principles of the two fundamental metabolic processes of the biosphere: photosynthesis and cellular respiration.

Highlighted in grey are the redox processes occurring on membrane systems that serve for energy conversion (photosynthesis, see 6.4; cellular respiration, see 6.10.3).

6.1.5. Energy Conversion and Energy Coupling

It follows from the Laws of Thermodynamics that the free energy change (∆G) of any series of coupled processes (e.g., chemical reactions) is equal to the sum of the Free Energy Changes of the individual reactions. This has profound implications for metabolism, as numerous endergonic metabolic processes can proceed spontaneously only when coupled with exergonic reactions such that the overall free energy change of the process is negative (∆G < 0), rendering the entire reaction exergonic. This phenomenon is termed energy coupling. It occurs primarily in sequential biochemical pathways and is a hallmark of virtually all metabolism. Specific energy-yielding reactions are repeatedly harnessed in metabolism to drive highly endergonic reactions. In cells, energy is most commonly stored in the form of adenosine triphosphate (Fig. 6.3; for structure, see Fig. 1.3).

Fig. 6.3. Energy coupling of exergonic and endergonic reactions in metabolic processes involving the adenylate system (ATP, ADP + P_i), illustrated by the coupling of phosphoenolpyruvate hydrolysis with the phosphorylation of glucose to glucose-6-phosphate

To a certain extent, other energy-rich nucleoside triphosphates can also be used in specific biosyntheses (e.g., of Nucleic Acids, see 1.2; carbohydrates, see 6.17.1; Lipids, see 6.11). The hydrolysis reaction ATP + H₂O -> ADP + P_i (P_i = inorganic phosphate) is a strongly exergonic process, as indicated by the standard molar Gibbs free energy (at pH 7) ∆G⁰' = -30.5 kJ·mol^-1. ATP formation according to the scheme: ADP + P_i

ATP + H₂O is strongly endergonic for these reasons: ∆G⁰' = +30.5 kJ·mol^-1. ATP can be formed by coupling with a suitable exergonic reaction. In this process, the phosphate group donor must possess at least an equivalent phosphate group energy, meaning its hydrolysis must yield a comparably high free energy change. This occurs, for example, during the hydrolysis of 1,3-bisphosphoglycerate or phosphoenolpyruvate (PEP) (see Table 6.3; Fig. 6.3). Such ATP-forming reactions are referred to as "substrate-level phosphorylation". However, the predominant share of ATP in plant cells is synthesized chemiosmotically, i.e., via energy coupling with a proton gradient generated in mitochondria during the oxidation of substrate molecules (see 6.10.3.3) and in chloroplasts during the light reactions of photosynthesis (see 6.4.9).¹

¹ In this context, the terms "membrane phosphorylation" or "membrane-level phosphorylation" are also used. — Ed. note

Thus, the second possibility for the energy Coupling of Endergonic and exergonic reactions involves utilizing the electrochemical energy of ion gradients (Fig. 6.4). In plants, these are gradients of hydrogen ions (protons) that exist across the Plasmalemma, tonoplast, and the membrane systems of Mitochondria and chloroplasts. In the latter two cases, they serve, as mentioned, for ATP synthesis; proton gradients across the plasmalemma and tonoplast are generated through the hydrolysis of ATP by proton-translocating ATPases (H⁺-ATPases, proton pumps). Additionally, a proton pump utilizing the energy of Pyrophosphate hydrolysis is present at the tonoplast. Transport processes, such as the translocation of hydrogen ions across a cell membrane mechanically coupled to the hydrolysis of an energy-rich bond (in this case, a phosphoanhydride bond), are termed primary Active Transport processes. As noted (see equation 6.19), the proton motive force driven by the gradient consists of two partial potentials: electrical and concentration-based. This electrochemical potential of hydrogen ions provides the driving force for the coupled (endergonic) transport of other ions or electroneutral metabolites (secondary active transport). This can utilize either the electrical component of the proton motive force (e.g., during ion transport through voltage-Gated Ion Channels — electrical coupling) or both the electrical and concentration proton gradients (electrochemical coupling), as occurs, for example, during the cotransport of hydrogen ions with electrically neutral metabolites via carrier Proteins (translocators). Ion channels generally conduct in both directions and are selective for individual ions or at least related types of ions; carriers are primarily highly selective for their specific substrate. Depending on the direction of transport, a distinction is made between uniporters, symporters, and antiporters (see Fig. 6.4). The most important currently known Primary and secondary active transport systems of the plasmalemma and tonoplast are summarized in Fig. 6.5.

Fig. 6.4. Energy coupling of exergonic and endergonic reactions via ion gradients across cell membranes (using the plasmalemma as an example).

In plants, this refers to hydrogen ion gradients. The endergonic accumulation of protons in the extracellular space is driven by the strongly exergonic hydrolysis of ATP. The resulting proton motive force is utilized for secondary active processes by ion channels (double blocks) or transporters (carriers, translocators) (circles). These are classified into uniporters, symporters, and antiporters depending on whether a single particle is transported (uniport) or two particles are transported in the same direction (symport) or in opposite directions (antiport). In principle, ion channels and transporters can also be involved in passive transport, which requires only a concentration gradient of the transported particle(s). Such protein-mediated transport processes that run solely until concentrations are equalized are termed Facilitated Diffusion. An example of a passive carrier is the triose phosphate/phosphate translocator of the inner chloroplast membrane (see Fig. 6.73). Porins of the outer membranes of chloroplasts and mitochondria, as well as peroxisomal (or glyoxysomal) membranes, also act as passive channels (see 6.5.6; 6.12).

Fig. 6.5. Examples of various primary active and passive or secondary active (shown in black) transport processes at the plasmalemma and tonoplast of plant cells. The stoichiometry of symporters and antiporters is not known in all cases; the figure merely illustrates the type of transported particle and the direction of transport. P-type ATPases form a phosphorylated intermediate during the transport cycle (P-phosphointermediate); V-type ATPases structurally resemble the ATP synthases of mitochondria (F₁/F₀-ATPase) or chloroplasts (CF₁/CF₀-ATPase) (V stands for vacuolar). ABC transporters utilize ATP energy to translocate larger organic compounds and complexes, such as the phytochelatin-Cd²⁺ complex (PC-Cd²⁺) or anthocyanin-Glutathione conjugates (anthocyanin-GS). ABC transporters are characterized by the presence of a specific Amino Acid Sequence required for ATP binding (ABC = ATP-binding cassette)

Fig. 6.6. Coupling of mechanical and chemical processes in motor proteins, illustrated by dynein-driven microtubule sliding (flagellar and ciliary motility, see 8.2).

Dynein is a very large protein complex with a molecular mass of 1–2 · 10⁶ Da, containing two to three force-generating heads (only one is shown) that form part of the heavy dynein subunit and possess ATPase activity. For clarity, only the heavy subunit of dynein is depicted, consisting of a base, HEAD, and stalk; the head and stalk together form the motor domain of dynein. In the axonemal structure of flagella, dynein binds to A-microtubules, while the rod-like stalks of the dynein heads tightly bind to specific sites on B-microtubules in the absence of ATP. The binding and hydrolysis of ATP in the dynein head induce a conformational change that is transmitted to the stalks. As a result, the stalk briefly detaches from the B-microtubule, moves toward the minus-end of the B-microtubule (toward the Base of the flagellum), and forms a new contact with a tubulin molecule. Upon the dissociation of ADP and P_i, the original conformation of dynein is restored, along with the bound state of the stalk on the B-microtubule. At this stage, mechanical force is transferred from dynein to the B-microtubule, causing it to slide relative to the A-microtubule toward the plus-end. Because the microtubules in the axoneme are linked by nexin bridges and anchored at the base, this sliding movement of microtubules results in the bending of the flagellum

The original function of the plasmalemma and tonoplast proton pumps was presumably to establish a cytoplasmic pH in plant cells within a narrow range between 7.5 and 8.0, maintaining disequilibrium primarily against the acidic extracellular environment and vacuolar contents. In marine algae, chloride is pumped into the cell by the action of a Cl^--translocating ATPase (chloride pump); by contrast, in salt glands (Limonium, Tamarix), Cl^- is extruded, with Na⁺ following via electrical coupling.

The Ca²⁺-translocating ATPases of the plasmalemma and Endoplasmic reticulum are responsible for removing Ca²⁺ that has passively diffused into the cell from the Cytoplasm, thereby maintaining a low cytosolic Ca²⁺ concentration (on the order of 10^-7 M).

The third possibility for energy coupling involves storing free energy in the form of an activated conformation of a protein molecule, the transition of which back to a lower-energy ground state is used to perform work. Motor proteins, such as dynein, convert the energy of the ATP phosphoanhydride bond into mechanical energy, where the phosphorylated form of the protein acts as the activated conformation (Fig. 6.6). The translocation of ions by ATPases is based on the differing Conformations of phosphorylated and unphosphorylated enzyme molecules. As a result, ion-binding sites with varying affinities are exposed on opposite sides of the membrane during the reaction cycle.

As has already become clear from several examples, energy coupling frequently occurs via energy conversion. For instance, plants first convert solar energy into electrochemical energy (electrical charge Separation and hydrogen ion concentration potential) and ultimately into chemical form (NADPH, ATP — see 6.4.4, Fig. 6.2). Motor proteins convert chemical energy into mechanical energy, ion-translocating ATPases convert chemical energy into an electrochemical potential, and the electrochemical potential of ion gradients (hydrogen ion gradients in plants) is utilized through A wide variety of transport processes to perform osmotic work (concentrating substances against an electrochemical potential gradient).

6.1.6. Enzymatic Catalysis

6.1.6.1. Basic principles of Catalysis

Equilibrium thermodynamics makes it possible to predict whether a given reaction is energetically feasible under specific conditions, i.e., whether it can proceed spontaneously. It also allows the calculation of reactant concentrations at equilibrium. However, equilibrium thermodynamics provides no information about the rate at which spontaneous reactions proceed. In reality, this rate can be extremely low. For example, the oxidation of glucose by oxygen is a highly exergonic process; yet, at physiological temperatures and normal pressure, glucose remains virtually indefinitely stable in the presence of oxygen. The primary reason for this is that the participants in chemical reactions must be in an activated

state from which the reaction proceeds. Bringing the reactants from their ground state into this energy-rich state requires an input of energy. The amount of energy required per mole of substance is termed the molar Gibbs energy of activation (∆G⁺), often referred to simply as the activation energy. In chemical reactions, activation can be achieved by raising the temperature. This increases the number of reactive molecules and thereby accelerates the reaction (typically doubling the rate for a 10 °C temperature rise). However, biochemical reactions must take place at relatively low temperatures. Moreover, living organisms experience only very minor temperature fluctuations, with metabolic processes occurring under virtually isothermal conditions. Raising the temperature as a means of accelerating metabolic reactions is therefore ruled out.

Catalysts (from Greek kata — down, lysis — loosening) are substances that, when added to a reaction mixture, increase the reaction rate while lowering the activation energy (Fig. 6.7). Catalysts emerge from the reaction unchanged and do not affect THE POSITION OF the reaction equilibrium (thus, they do not alter the standard Gibbs free energy of the reaction, ∆G). They merely accelerate the reaction and, consequently, the attainment of equilibrium. Because they are neither altered nor consumed during the reaction, they can be used over and over again to perform their function and need only be present in small amounts.

Fig. 6.7. Free energy profile of a system over the course of uncatalyzed and catalyzed reactions, using the disproportionation of H₂O₂ as an example.

For the enzyme-catalyzed reaction (shown by the gray line), the reaction pathway is depicted in greater detail. The lengths of the arrows in the graph are proportional to the corresponding free energy changes of the H₂O₂ disproportionation reaction; E — enzyme; S — substrate; ES — enzyme-substrate complex; EP — enzyme-product complex; P — reaction product

Accelerating metabolic reactions is The primary function of biocatalysts. With the exception of a few catalytically active Ribonucleic Acids (ribozymes), biocatalysts are proteins. They are referred to as enzymes (from Greek zymē — leaven) or ferments (from Latin fermentum — leaven) and follow the same fundamental principles as chemical catalysts; enzymes likewise lower the activation energy of the catalyzed reaction (Fig. 6.7) without altering its equilibrium (and thus ∆G).

For instance, the molar standard Gibbs free energy of activation (∆G⁺ for the disproportionation reaction of hydrogen peroxide H₂O₂ -> H₂O + 1/2O₂ is: ∆G⁺ = +75 kJ • mol^-1; this energy barrier can be overcome by heating the H₂O₂ solution. In the presence of finely divided platinum, ∆G⁺ = +49 kJ • mol^-1. Platinum acts as a catalyst, allowing the reaction to proceed at a measurable rate even at room temperature. The enzyme catalase catalyzes this disproportionation with an activation energy of ∆G⁺ = +23 kJ • mol^-1. With catalase present, hydrogen peroxide rapidly decomposes into H₂O + 1/2O₂ at room temperature. In each case, the molar standard enthalpy of this exergonic reaction is ∆G⁺ = -97 kJ • mol^-1. Enzymes are exceptionally efficient catalysts. For example, Carbonic anhydrase accelerates the Hydration of CO₂ according to the equation H₂O + CO₂ ⇄ H₂CO₃ by a factor of 10⁷, achieving a turnover rate of 10⁵ CO₂ molecules per second per enzyme molecule.

An enzymatic reaction (see Fig. 6.7) begins with The formation of an enzyme-substrate complex (ES); this complex then converts into an enzyme-product complex (EP), which rapidly dissociates to release the reaction products and regenerate the free enzyme. Typically, the overall activation energy required is determined by the energy barrier that must be surmounted to convert ES into EP. This step is rate-limiting for the entire reaction.

6.1.6.2. Molecular Mechanisms of Enzymatic Catalysis

Enzymes exhibit both substrate Specificity and reaction specificity. The degree of substrate specificity varies among individual enzymes. Some Hydrolases, for instance, are relatively non-specific and hydrolyze a wide range of substrates, whereas others show specificity for particular molecular groupings. Thus, α-glucosidases hydrolyze α-glucosidic bonds in various substrates, but not β-glucosidic bonds. Many enzymes display an exceptionally high degree of substrate specificity. A striking feature is their often remarkable ability to recognize stereoisomers. This refers to the drastically different turnover rates of metabolites that differ only in the spatial arrangement of their substituent groups (such as cis-trans isomers or optical isomers, which relate to each other as an object and its mirror image — known as enantiomers).

Substrate specificity is based on a precise, complementary fit between the substrate and the catalytically Active Site of the enzyme, known as the active center. In the simplest case, the substrate and the catalytic site fit together like a lock and key. However, this metaphor, introduced by Emil Fischer as early as 1890, fails to account for the fact that enzyme-substrate binding is often a dynamic process accompanied by Conformational Changes in both the enzyme and the substrate. This concept, postulated by D. E. Koshland in 1958, is known as the induced fit model. The active center of an enzyme is frequently formed only after substrate binding and the resulting conformational change take place, as seen in phosphoglycerate kinase (Fig. 6.8).

Fig. 6.8. Substrate-induced conformational change (induced fit), illustrated schematically using phosphoglycerate kinase. Upon binding of the substrates ADP and 1,3-bisphosphoglycerate, the conformation of the protein undergoes a profound alteration, causing both domains of the enzyme to close over the bound substrates while simultaneously excluding water molecules from the interior (center). Phosphatryl transfer takes place within this newly formed, water-free microenvironment. Following the relaxation back to the 'open' conformation, the reaction products diffuse away from the catalytic site. A schematic cross-section through the active center of the enzyme is shown, with reaction components depicted at approximately equal scale.

Table 6.4. International Classification of Enzymes: class designation, code number, and type of catalyzed reaction

Enzymes are generally named after the experimentally discovered reaction they catalyze, although under physiological conditions within the cell they can also catalyze the reverse reaction (example: shikimate dehydrogenase, see 6.13.2, Fig. 6.107). The classification follows rules established by the Enzyme Commission (EC) of the IUB (International Union of Biochemistry). Each enzyme is assigned a 4-part code number: for example, E.C. 1.1.1.25 is the code for shikimate dehydrogenase (see Fig. 6.107):

1. Oxidoreductases (oxidation-reduction reactions)

1.1. Acting on the > CH—OH group

1.2. Acting on the > C—O group

1.3. Acting on the > CH—CH< group

1.4. Acting on the > CH—NH₂ group

1.5. Acting on the > CH—NH— group

1.6. Acting on NADH or NADPH

2. Transferases (transfer of functional groups)

2.1. C₁ groups

2.2. Aldehyde or ketonic groups

2.3. Acyl groups

2.4. Glycosyl groups

2.5. Alkyl or aryl groups (except methyl groups)

2.6. N-containing groups

2.7. P-containing groups

2.8. S-containing groups

3. Hydrolases (hydrolytic reactions)

3.1. Ester bonds

3.2. Glycosidic bonds

3.3. Ether bonds

3.4. Peptide bonds

3.5. Other C—N bonds

3.6. Acid anhydride bonds

4. Lyases (cleaving C—C, C—O, C—N, and other bonds)

5. Isomerases (isomerization, i.e., intramolecular rearrangements)

5.1. Racemases and epimerases

5.2. Cis-trans-isomerases

5.3. Intramolecular oxidoreductases

6. Ligases (synthetases*) (covalent joining of two molecules coupled with ATP hydrolysis)

* Enzymes catalyzing anabolic reactions that proceed without ATP cleavage are referred to as synthases.

Phosphoglycerate kinase, a glycolytic enzyme (see 6.10.1), binds 1,3-bisphosphoglycerate and adenosine diphosphate (ADP) and catalyzes The transfer of a phosphoryl group from the carboxyl group of 1,3-bisphosphoglycerate to ADP, yielding ATP and 3-phosphoglycerate. This process involves the cleavage of one anhydride bond (in 1,3-bisphosphoglycerate) and the formation of another (in ATP). Such a reaction would be entirely impossible in an aqueous environment, as hydrolysis is thermodynamically favored. The solution to this problem lies in the fact that the binding of ADP and 1,3-bisphosphoglycerate induces a conformational fit in which both domains of the enzyme (see Fig. 6.8) close over the bound substrates, excluding water. Only in this state is the active site formed and the Phosphate group transfer enabled. Upon completion of catalysis, the "open" conformation is restored, and the reaction products dissociate from the enzyme.

In addition to substrate specificity, enzymes exhibit reaction specificity. This means that a biocatalyst catalyzes only one of many thermodynamically possible pathways for substrate conversion. Regarding the underlying mechanisms, there are relatively few reaction types, which form The basis of the systematic enzyme nomenclature (Table 6.4).

The name of an enzyme (for substrate-cleaving enzymes) is typically formed by adding the suffix -ase to the substrate name: for example, proteinase for protein-cleaving, amylase for starch-hydrolyzing (from Latin amylum), and lipase for fat-cleaving (from Greek lipos) enzymes. In addition, traditional historical names remain in common use, such as Pepsin or catalase. A unified, systematic, and mandatory international Classification and Nomenclature for all known enzymes was proposed by the International Enzyme Commission, assigning each enzyme a unique classification number (EC number) for unambiguous identification (see Table 6.4). Because systematic names can be quite cumbersome, shorter trivial names continue to be widely used alongside them.

While for some enzymes the protein moiety alone is catalytically active, others require additional substances (Cofactors). These cofactors can be Metal Ions (e.g., Mg²⁺, Mn²⁺, Zn²⁺, Fe²⁺, Fe³⁺, Cu²⁺, K⁺), which may be required to anchor the substrate to the enzyme molecule or participate directly in the catalytic reaction as part of the active site. When organic compounds are required as cofactors,

they are termed Coenzymes. If a coenzyme is bound to the protein moiety so tightly that it is difficult to dissociate (e.g., it cannot be removed from the enzyme complex by dialysis), it is referred to as a prosthetic group (from Greek prosthetos, added). For example, in Cytochromes, the heme group is covalently bound to the protein (see 6.4.6; 6.10.3.3). The complete enzyme-cofactor complex is called the holoenzyme, whereas the catalytically inactive protein component alone is termed the apoenzyme.

If cofactors are consumed stoichiometrically relative to the substrate (as is typical in redox reactions), they can also be designated as cosubstrates.

6.1.6.3. Kinetics

The enzyme-catalyzed reaction converting a substrate into its product proceeds according to the general scheme shown in Fig. 6.7. To introduce the Basic Concepts of enzyme kinetics, the reaction can be represented in a simplified form

Here, k₊₁, k_-1, and k₊₂ denote the rate constants of the individual partial reactions (in reciprocal seconds). In the simplified model, we assume that the reverse reaction E + P —> ES proceeds negligibly slowly (k₂ ≈ 0), and the dissociation of the enzyme-substrate complex (ES) into enzyme + product (P) is much slower than the reverse reaction ES —> E + S (k₊₂ ⋘ k_-1). Therefore, the rate-determining step of the entire reaction is the conversion ES —> E + P, since the slowest partial reaction dictates the overall reaction rate. Under these conditions, The rate of substrate conversion into its product is given by the equation

The maximum velocity is reached when all of the enzyme is present in the form of the enzyme-substrate complex:

At a velocity equal to half of the maximum velocity (1/2V_max), the concentration of free enzyme equals that of the enzyme-substrate complex ES: [ES] = [E]. At equilibrium, the Formation of the enzyme-substrate complex is described by the expression

and thus, after rearrangement, assuming that k₊₂ ⋘ k_-1

The ratio k_-1/k₊₁ is called the Michaelis constant (K_m). It can be defined as the Substrate Concentration at which [ES] = [E] (see above). Thus, K_m indicates the substrate concentration at which exactly half of the maximum reaction velocity is achieved.

Since neither V_max (and thus 1/2V_max) nor the corresponding substrate concentration can be determined with sufficient accuracy from the plot of v versus [S] (Fig. 6.9, A), K_m is best determined after linear transformation of the graph shown in Fig. 6.9, A, which is achieved by plotting the same relationship using reciprocal coordinates (1/v versus 1/[S] — the Lineweaver–Burk plot; Fig. 6.9, B). For a given enzyme, a given substrate, and a specific temperature, K_m is a constant expressed in mol L^-1. K_m values can also be determined for cofactors.

Fig. 6.9. Dependence (A) of the reaction velocity (v) on the substrate concentration [S] for an enzyme-catalyzed reaction according to the Michaelis–Menten model (see equations 6.25 and 6.26) (after Lineweaver–Burk). B — v_max and K_m can be determined more accurately by plotting the data in reciprocal coordinates

6.1.6.4. Environmental Effects on Enzyme Activity

Enzyme activity is largely determined by temperature, pH, and the ionic COMPOSITION OF THE medium. These factors affect The structure of the enzyme protein. The dependence of the reaction rate on temperature follows a curve with an optimum (Fig. 6.10). The temperature optimum varies among individual enzymes and frequently lies between 30 and 50 °C. Before reaching the optimum, the reaction rate doubles or triples with each 10 °C rise in temperature. The ratio of reaction rates v_T+10/v_T is called the Q₁₀ coefficient. For enzyme-catalyzed reactions, Q₁₀ = 2–3. At temperatures above the optimum, activity typically drops very rapidly, which is explained by the thermal Denaturation of the enzyme protein; due to a sharp increase in entropy, denaturation becomes the favored pathway. Certain enzymes are remarkably thermostable. For example, Ribonuclease and peroxidase can withstand even boiling. Freezing is tolerated by most enzymes without damage; therefore, enzyme solutions are usually stored in a frozen state.

Fig. 6.10. Dependence of the velocity (v) on temperature for a non-catalyzed (or non-protein-catalyzed) and an enzyme-catalyzed chemical reaction. The temperature optima of most enzymes lie between 30 and 50 °C

Ionizable groups of the substrate or the enzyme are frequently involved in binding, in the catalytic Conversion of the substrate, or in establishing the conformation of the enzyme protein¹. Consequently, enzyme activity depends on the ambient pH. This dependence can be very pronounced, and the pH optima of different enzymes—or of the same enzyme with different substrates—may lie at quite distinct pH values.

¹ Proteins typically contain ionizable groups capable of reversibly binding H⁺ ions: —COOH, —NH₂, etc. The degree of ionization depends on the H⁺ concentration in solution, i.e., on the pH. — Ed. note.

For instance, Plasma Membrane H⁺-ATPase has a pH optimum of 6.5, arginase for Arginine has a pH optimum of 9.7; fumarase has two pH optima: 6.5 with fumarate as a substrate, and 8.5 with malate; acid Phosphatases exhibit a pH optimum around 5. Over a wide pH range, activity remains unchanged—for example, in invertase, which cleaves the electrically neutral substrate sucrose and for which both extracellular and intracellular isozymes are known (see 6.8.4)

Since numerous cellular enzymes have different pH optima and pH values vary across individual compartments, intracellular pH shifts significantly impact metabolism. The ionic potential ("Ionic strength") can also affect enzyme proteins. Among other things, the ionic potential influences enzyme conformation via their degree of hydration.

Ultimately, enzyme activity also depends directly on the substrate concentration and, if the enzyme functions in conjunction with a dissociable cofactor, on the cofactor concentration (see Fig. 6.9). At metabolic branch points, the concentration of the shared substrate can determine which pathway is favored. When substrate is limited, the enzyme with the lower Michaelis constant will preferentially operate.

In organisms capable of Alcoholic Fermentation (such as Yeasts), pyruvate can either be decarboxylated by pyruvate decarboxylase to form acetaldehyde, or undergo oxidative decarboxylation through the action of pyruvate dehydrogenase. At low pyruvate concentrations, because of the lower Michaelis constant of pyruvate dehydrogenase, the Formation of Acetyl-CoA proceeds preferentially, whereas at higher concentrations of pyruvate, acetate formation comes to the forefront (Fig. 6.11).

Fig. 6.11. The direction of metabolite fluxes at a metabolic branch point depends on the concentration of the shared substrate and the K_m values of the competing enzymes at the junction. The lower the substrate concentration, the more preferentially the reaction proceeds via the pathway catalyzed by the enzyme with the lowest K_m value. Pyruvate metabolism via pyruvate dehydrogenase or pyruvate decarboxylase is shown as an example

6.1.7. Introduction/15.html">Regulation of enzyme Activity

Like all proteins, enzymes are constantly synthesized and degraded within the cell. The rate of a given cellular reaction is primarily controlled by the amount of enzyme present. While this process is crucial for adapting to altered metabolic demands, it is far too slow to fine-tune metabolism. Therefore, alongside controlling enzyme quantities, numerous efficient and predominantly reversible processes exist to directly regulate enzyme activity, enabling the cell to rapidly and flexibly adjust its metabolism to changing requirements.

6.1.7.1. Control of Enzyme Abundance

The amount of a specific protein in a cell is the net result of its rates of Synthesis and degradation. Protein Synthesis depends on the transcriptional activity of the encoding Gene (see 7.2.2) and post-transcriptional processes that take effect following mRNA synthesis. The latter determine, for instance, mRNA stability and, consequently, its cellular abundance alongside Transcription rates. Further regulatory mechanisms involve Translation—influencing the Translation of the nucleic acid code into the corresponding linear amino acid sequence (see 7.3.1.2)—and, where necessary, the Processing of the primary polypeptide into a mature, enzymatically active protein. For example, the lipid-cleaving lipase in germinating Ricinus seeds is released from a precursor protein through the action of a proteinase.

In many cases, plants contain several isozymes. Isozymes are enzymes that catalyze the same chemical reaction but

differ in their chemical properties (such as their isoelectric point—see 1.3.1—or pH optimum—see 6.1.6.4). Isozymes are frequently products of different genes, though they can also arise from post-transcriptional modifications leading to distinct variant forms. In enzymes with a quaternary structure (see 1.3.2.3), the number of isozymes can be further increased through complex formation between protomer isoforms (heterooligomerization). Gene families encoding isozymes offer the advantage that each gene, driven by an individual promoter, can specifically regulate its own transcription (see 7.2.2.3). This allows the organism to maintain a targeted distribution of enzymatic activity—depending on the cell compartment, tissue, or developmental stage—or to respond to a variety of natural stimuli. For instance, when metabolic demand rises, an inducible enzyme can be synthesized in Addition to a constitutively active enzyme. Finally, isozymes may differ in the mechanisms that control their activity (see Fig. 6.14).

Control of enzyme abundance is less pronounced in primary metabolism than in enzymes performing specialized functions, or those synthesized only upon specific demand or produced in increased amounts. Examples include plant defense responses against pests or pathogens (see Chapter 9). One primary metabolic enzyme regulated, among other ways, at the level of abundance is nitrate reductase. The synthesis of this enzyme is induced by nitrate (NO^-₃) and repressed by ammonium (NH⁺₄). In contrast, constitutive enzymes, which are responsible for maintaining basic cellular functions (housekeeping enzymes), are predominantly regulated through activity control mechanisms.

6.1.7.2. Control of Enzyme Activity

Enzymes can alter their activity through reversible covalent modification or non-covalent interactions with regulatory molecules (modulators). Frequent covalent modifications include phosphorylation and dephosphorylation, catalyzed by specific protein Kinases or phosphoprotein phosphatases. Typically, ATP serves as the phosphate donor (Fig. 6.12, A), while Serine, Threonine, Tyrosine, or Histidine act as acceptor Amino Acids in regulatory phosphorylation. For instance, phosphoenolpyruvate carboxylase is phosphorylated at a specific serine residue and thereby activated; pyruvate orthophosphate dikinase is inactivated via phosphorylation at a specific threonine residue (see 6.5.8, 6.5.9). Numerous enzymes, such as those located in chloroplasts and mitochondria, undergo redox control via dithiol-disulfide modification (see Fig. 6.71). Thioredoxins—small proteins with a molecular mass of about 12 kDa that occur in multiple isoforms within the cytoplasm, mitochondria, and Plastids of plants—supply numerous reducing equivalents in the process. Examples of enzymes regulated via dithiol-disulfide modification by thioredoxin include enzymes of The Calvin Cycle, notably fructose-1,6-bisphosphatase and phosphoribulokinase (see 6.5.3). Because photosynthetic electron transport operates only in the light and generates reduced thioredoxin, the redox control of Calvin cycle enzymes serves to adapt photosynthetic CO₂ fixation activity to the alternation of day and night (see 6.5.5).

Fig. 6.12. Regulation of enzyme activity through reversible covalent modifications: A — phosphorylation-dephosphorylation; B — regulation of E. coli Glutamine Synthetase by adenylylation-deadenylylation

In Escherichia coli, glutamine synthetase can exist in an active or inactive form; in the latter state, the 12 subunits of the enzyme are bound to 12 adenosine monophosphate (AMP) molecules. The enzyme catalyzing the adenylylation reaction (which inactivates glutamine synthetase) is inhibited by 2-oxoglutarate and activated by glutamine. Conversely, the deadenylylating enzyme, which restores glutamine synthetase to its active form, is inhibited by glutamine and stimulated by 2-oxoglutarate (Fig. 6.12, B). This system establishes an automated self-regulation of glutamine synthesis from 2-oxoglutarate: if abundant 2-oxoglutarate and little glutamine are present, the synthesis of this amino acid increases; conversely, it halts.

Alterations in enzyme activity via non-covalent interactions can take place either at the catalytic site itself or at a distant Location. If a molecule structurally related to the substrate binds to the catalytic site but cannot be metabolized, this is referred to as competitive inhibition, since an excess of the substrate can displace the inhibitor. The extent of competitive inhibition thus depends on the concentration ratio between the inhibitor and the substrate. Kinetically, a competitive inhibitor is characterized by the fact that V_max remains unchanged in the presence of the inhibitor, whereas the K_m value increases. When the product of an enzymatic reaction acts as a competitive inhibitor to the substrate of that same reaction, it is called product inhibition. This mechanism ensures that only as much substrate is consumed as can be processed in subsequent reactions, thereby preventing the accumulation of unneeded metabolic intermediates.

Competitive Inhibitors (Fig. 6.13) can be exceptionally potent. This is especially true when they act as structural analogs of the Transition State of the activated substrate. In such cases, even a large excess of substrate makes them difficult to displace from the enzyme. Sulfonamides are classic examples of competitive inhibitors; these antibacterial agents act by competitively blocking the incorporation of structurally similar Para-aminobenzoic Acid into Folic acid—an essential compound required for the Synthesis of purine NUCLEOTIDES (see 6.14). Because humans do not synthesize folic acid themselves but instead obtain it as a vitamin from their diet, sulfonamides do not inhibit human metabolism ¹

¹ It follows that administering sulfonamide drugs concurrently with Vitamins (folic acid) is pointless, as the bacteria would not be killed in this case — Ed. note.

Fig. 6.13. Examples of competitive inhibitors and their corresponding enzyme substrates (see text)

Monofluoroacetic acid (CH₂F—COOH) is a toxic substance synthesized in the leaves of Dichapetalum cymosum (family Dichapetalaceae), a South African plant highly toxic to grazing livestock. Monofluoroacetate can bind to coenzyme A in place of an acetyl group and subsequently be transferred by the enzyme citrate synthase onto oxaloacetate (for The Citric Acid Cycle, see 6.10.3.2), yielding monofluorocitrate. This compound is a highly potent competitive inhibitor of aconitase, the enzyme that processes citrate within The Citric Acid cycle. In the plant Dichapetalum, the corresponding toxic effect is presumably prevented because the toxin does not reach its specific Site of Action (i.e., the mitochondria) but remains sequestered in a separate compartment—the vacuole.

Fig. 6.14. Fine-tuning of parallel metabolic pathways via negative feedback inhibition of allosterically regulated isozymes by end products. The accumulation of product Z inhibits its own formation without affecting the metabolic pathways leading to products X and Y, which branch off from the common intermediate B

In allosterically regulated enzymes, the binding of a modulator induces a conformational change in the enzyme (from Greek allos = other, stereos = shape), which either inactivates the catalytic center (the modulator acts as an allosteric inhibitor) or activates it (the modulator acts as an allosteric activator). If the modulator is identical to the substrate, the enzyme is termed a homotropic enzyme; if it differs from the substrate, it is a heterotropic enzyme. Allosteric control is widespread in metabolism and highly efficient. Key metabolic enzymes—typically those catalyzing the first step of a reaction chain and inhibited by the accumulating end product of the entire pathway—are frequently regulated in this manner. This negative feedback (feedback inhibition) is extremely economical, as it ensures that metabolite flux through complex metabolic pathways is adjusted precisely to physiological demand. If the intracellular concentration of the end product drops, the allosteric inhibitor dissociates from the enzyme, and substrate processing resumes or becomes activated. Due to the presence of isozymes in branched reaction chains, individual branches can be regulated independently by their respective end products, which inhibit the corresponding isozyme (the first one following the branch point) via negative feedback (Fig. 6.14; see also Figs. 6.108 and 6.112). Positive feedback occurs when a modulator activates an allosteric enzyme.

Allosterically regulated enzymes typically consist of multiple subunits whose activities are interdependent. This property is known as cooperativity. The conformational change triggered by modulator binding (in homotropic enzymes, substrate binding at one of the catalytic centers; in heterotropic enzymes, regulator binding at a different site on the complex) is transmitted to the other subunits, altering (usually increasing) the affinity of the remaining catalytic centers for the substrate. Consequently, substrate saturation curves for allosterically regulated enzymes exhibit a sigmoidal shape (Fig. 6.15). Accordingly, these enzymes process substrate efficiently only above a specific threshold concentration, beyond which even minor shifts in substrate concentration lead to drastic changes in conversion rates.

Fig. 6.15. Effect of Substrate concentration [S] on the reaction velocity [V] catalyzed by a homotropic, allosteric enzyme. Allosteric enzymes often consist of multiple subunits (two are shown) and exist in a low-activity form (squares) in the absence of substrate. Binding of the substrate to one subunit induces a transition of all subunits into a high-activity form (circles). As a result, at low substrate concentrations, the reaction velocity increases slowly at first and then exponentially. Once the enzyme has transitioned into the activated form, the dependence of reaction velocity on [S] approaches Michaelis–Menten kinetics

6.1.7.3. Regulation through the Assembly of Enzymes into Multienzyme Complexes or Compartments

A crucial foundation for the orderly and controlled progression of metabolism within the cell is the association of enzymes for specific reaction sequences (reaction cycles) into multienzyme complexes or entire metabolic regions within specific compartments.

In a multienzyme complex, several individual enzymes are assembled into a higher-order structure. This organization ensures the rapid, coordinated conversion of a substance through multiple sequential steps. When intermediate products cannot be detected externally, this phenomenon is referred to as metabolite channeling. Examples of multienzyme complexes include the pyruvate dehydrogenase complex (see 6.10.3.1) or the Yeast fatty acid synthase complex (see 6.11.1). A specific enzyme complex in filamentous fungi combines five enzymatic activities of the aromatic Amino acid Biosynthesis pathway—which are handled by separate enzymes in Escherichia coli—into a single, multifunctional polypeptide (see 6.13.2).

The integration of entire groups of enzymes, cofactors, and metabolites within reaction spaces separated from their surroundings by metabolic barriers (i.e., compartments such as the cytoplasm, chloroplasts, and mitochondria) is critical for the orderly functioning and Regulation of cellular metabolism. Metabolite exchange between compartments occurs primarily via specific carriers, or transporters (translocators) (see Fig. 6.4), the activity of which is also subject to various regulatory mechanisms.

The principles of bioenergetics, enzymatic catalysis, and regulation discussed in this section will provide a useful framework for understanding the diverse plant functions presented in the subsequent chapters.