lecture 21 - introduction to metabolism:...

Bioc 460 - Dr. Miesfeld Fall 2008

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Figure 1.

Solar energy

Lecture 21 - Introduction to Metabolism: Bioenergetics Key Concepts

• Energy conversion in biological systems • Metabolic redox reactions • Review of thermodynamic principles and coupled reactions • The adenylate system and the Energy Charge of the cell

KEY CONCEPT QUESTIONS: How is energy from the sun converted to chemical energy? What is reaction coupling and why is it important in metabolic pathways? ENERGY CONVERSION IN BIOLOGICAL SYSTEMS Essentially all biological processes on this planet are directly or indirectly affected by oxidation-reduction reactions in photosynthetic organisms that convert solar energy into chemical energy. Photosynthetic organisms use this chemical energy to sustain life during the daylight hours and to produce carbohydrates from CO2 that can be stored as metabolic fuel for use at night. All other organisms obtain chemical energy from their environment, which in many cases, means consuming the organic materials produced by photosynthetic organisms and using them as metabolic fuel for aerobic respiration. Bioenergetics is a term that describes the processes involved in these energy conversion reactions in living systems. One branch of biochemistry is devoted to understanding the molecular mechanisms that control these processes.

Living organisms need an constant input of energy to put off death as long as possible. The reason for this is that life is dependent on the maintenance of a highly ordered steady state called homeostasis which requires energy. An organism that is at equilibrium with its environment is no longer alive, indeed, organisms must maintain a steady state that is far from equilibrium in order to survive. For example, the concentration of glucose needed to sustain life in a saguaro cactus is much higher inside the cactus than it is in the surrounding desert, and it requires the process of photosynthesis to provide the energy needed to keep it this way (figure 1). Similarly, the concentration of sodium chloride is lower inside the cells of a humpback whale than it is in the surrounding ocean. In this case, it is the whale's diet of shrimp and plankton that provide the chemical energy required to maintain a safe intracellular sodium chloride concentration. When an organism can no longer maintain homeostasis using these energy conversion processes, the intracellular concentration of water, essential ions, and macromolecules, begin to equilibrate with the surroundings and the organism dies. The reason all living organisms need an input of energy (which can be stored), is to delay reaching equilibrium with the environment as long as possible.

Energy conversion in living systems is required for three types of work that maintain the steady state, 1) chemical work in the form of macromolecular biosynthesis of organic molecules, 2) osmotic work to maintain a concentration of intracellular salts and organic molecules that is different than the extracellular milieu, and 3) mechanical work in the form of flagellar rotation or


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Figure 2.

OIL RIG

Figure 3.

muscle contraction. The conversion of hydrogen to helium by thermonuclear reactions in the sun, and the subsequent release of energy in the form of visible light, is called solar energy and it is the ultimate power source for life on earth (figure 2). Solar energy provides all of the energy required for the two types of organisms that inhabit this planet, photosynthetic autotrophs and heterotrophs. Photosynthetic autotrophs use solar energy to oxidize H2O and generate chemical energy that is used to maintain homeostasis during daylight hours. Photosynthetic autotrophs also use this chemical energy to convert atmospheric CO2 into carbohydrate (C6H12O6) which is a storage form of energy used at night. The process of oxidizing H2O to capture chemical energy and generate O2 is called photosynthesis, whereas, the conversion of CO2 to C6H12O6 is carbon fixation. The most abundant photosynthetic autotrophs in the biosphere are vascular plants, single cell algae, and photosynthetic bacteria.

Heterotrophs, which includes all non-photosynthetic organisms, are dependent in one way or another on photosynthetic autotrophs as a source of chemical energy (carbohydrate) which is used as metabolic fuel for aerobic respiration. Importantly, the production of O2 by photosynthetic autotrophs through the oxidation of H2O is critical for aerobic respiration because O2 is the terminal electron acceptor in this process. Some bacteria derive redox energy from compounds in the soil and are considered heterotrophs, although they are not reliant on photosynthetic autotrophs. Metabolic redox reactions Both photosynthesis and aerobic respiration interconvert energy using a series of linked oxidation-reduction reactions in which electrons are transferred from a molecule of higher electrochemical potential to one of lower electrochemical potential. Oxidation Is the Loss of electrons and Reduction Is the Gain of electrons (OIL RIG). A series of linked oxidation-reduction reactions, often called redox reactions, transfer electrons from one compound to another in sequential fashion. Since electrons do not exist free in solution, a reduced compound becomes oxidized when it transfers an electron to an oxidized compound that becomes reduced (figure 3). The importance of redox reactions in biochemical processes is that chemical work can be performed using the energy made available by electron transfer.

The initiating biochemical event required for all subsequent energy conversion processes in our biosphere is the absorption of light energy by pigment molecules, such as chlorophyll, present in photosynthetic organisms. Light absorption causes photooxidation of chlorophyll which results in the transfer of an electron from the chlorophyll molecule to an acceptor molecule which then passes the


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Figure 4.

Figure 5.

electron to another acceptor molecule of lower electrochemical potential. The redox reactions of aerobic respiration are fundamentally similar to that of photosynthesis, except in this case, oxidation of glucose by enzyme-mediated reactions transfers 24 e- to the coenzymes nicotinamide adenine dinucleotide (NAD+) and flavin adenine dinucleotide (FAD). Oxidation of NADH and FADH2 provides redox energy for the process of oxidative phosphorylation which synthesizes ~32 ATP for each glucose molecule oxidized (figure 4). The electrons are passed through electron carrier molecules associated with the mitochondrial inner membrane and are ultimately donated to O2 as the final electron acceptor to produce 12 H2O. The biochemistry of aerobic respiration (glycolysis and the citrate cycle), and oxidative phosphorylation (electron transport system and ATP synthesis), are described in later lectures. Review of thermodynamic principles To better understand how energy is utilized in biological systems, we need to review basic thermodynamic principles in the context of a system and its surroundings. The system defines the collection of matter in a defined space, whereas, the surroundings is everything else. The system and surroundings together constitute the universe. Biological systems are open systems as both matter (nutrients and waste products) and energy (primarily heat) are freely exchanged with the surroundings. Three laws of thermodynamics have been defined. The first law of thermodynamics states that in any physical or chemical change that occurs, the total amount of energy in the universe remains the same, even though the form of energy can change. Put another way, energy can neither be created or destroyed, only transformed. The second law of thermodynamics states that in the absence of an energy input, all natural processes in the universe tend toward disorder (randomness), and moreover, that the measure of this disorder, called entropy, is always increasing in the universe. The Gibbs free energy value is the third bioenergetic principle we need to review. This is usually expressed as a change in free energy and relates to the spontaneity of a chemical reaction. In fact, all three of these terms can be tied together and are related to the equilibrium constant (Keq) of a reaction. First Law of Thermodynamics The first law of thermodynamics states that energy cannot be created or destroyed, only converted from one form to another. In a closed system such as a bomb calorimeter in which a compound is combusted by a spark under constant pressure in the presence of pure oxygen (completely oxidized), the amount of heat exchange between the reaction chamber and a surrounding water jacket is a measure of the


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Figure 6.

Figure 7.

change of enthalpy, ΔH (figure 5). A reaction that gives off heat is called exothermic and ΔH has a negative value, whereas, a reaction that absorbs heat is called endothermic and ΔH is a positive number. The enthalpy in a chemical compound is reflected in the number and type of chemical bonds. Therefore, depending on how a chemical reaction changes the number and type of bonds in the reactants and product, the reaction will either be exothermic or endothermic. For example, the combustion of 1 gram of glucose is an exothermic reaction and produces heat, CO2, and H2O. The temperature increase of the surrounding water is a measurement of the potential energy of the glucose as reflected in the number and type of chemical bonds. Complete oxidation occurs by igniting a spark in the presence of pure O2, which in this example, converts glucose (C6H12O6) to CO2 and H2O in an exothermic reaction that raises the temperature of the surrounding water +3.75ºC. We can write an equation for this reaction in terms of chemical reactants and products in which the heat produced is at constant pressure:

C6H12O6 + 6 O2 6 CO2 + 6 H2O + Heat A Calorie (C) is a unit of energy that was originally defined by the amount of heat energy required to raise 1 kilogram of water from 14.5 ºC to 15.5 ºC using a calorimetry device. Energy can also be expressed in the international unit of measurement the Joule (J) in which 1 Calorie = 4.184 kJ. Note that in nutritional sciences calorie with a capital "C" actually refers to kilocalories (kcal), so 1 kcal = 4.184 kJ. For example, as shown in figure 6, the energy potential of 1 gram of glucose is 15.7 kJ regardless of the route taken to fully oxidize it (bomb calorimeter or a mouse). Second Law of Thermodynamics The second law of thermodynamics states that all natural processes in the universe tend towards disorder (randomness) in the absence of energy input. This concept of disorder is defined by entropy (S), a measure of disorder. The more disorder in a system, the higher the value of entropy. The second law of thermodynamics is used in biochemistry to determine the directionality of a reaction, that is, its spontaneity. Irreversible reactions, such as an ice cube melting at room temperature, is an example of an increase in entropy, ΔSuniverse > 0. The H2O molecules in the ice


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crystals are highly ordered through hydrogen bonding, whereas the H2O molecules in liquid water are much more disordered (many freedoms of rotation). Ice melting at room temperature is spontaneous, however, the process can be reversed by the input of energy, namely the electricity used by a freezer compartment to lower the temperature of an ice cube tray (figure 7). Once the water is frozen solid in the freezer compartment, the input of electricity restrains the ice crystals from melting again. Similarly, the metabolic energy required for sustaining life restrains the natural tendency of the molecules within the organism to become disordered as dictated by the second law of thermodynamics. Two examples of increased entropy in living systems are the degradation of biomolecules into a larger number of smaller components, and random mutations in DNA that result in disorganized genetic information. Gibbs Free Energy (G) and the Equilibrium Constant (Keq) In 1878, an American theoretical physicist name J. Willard Gibbs, described a way to determine if a chemical reaction is favorable or unfavorable under constant pressure and temperature using a function he called free energy. The Gibbs free energy term (G), is defined as the difference between the enthalpy of the system (H) and the entropy (S) at a given temperature (T). Since the absolute values for G, H, and S cannot be easily determined, but the change between two states can, the most useful Gibbs free energy equation is:

ΔG = ΔH - TΔS Importantly, if the ΔG value for a reaction is less than zero (ΔG<0), then the reaction is favorable and is exergonic, however if the ΔG value is greater than zero (ΔG>0), then the reaction is unfavorable and endergonic. A reaction in which ΔG is equal to zero (ΔG=0) is at equilibrium, meaning that the rate of formation of products is equal to the rate of formation of reactants (no net product or reactant formation is occurring). This relationship between free energy, enthalpy, and entropy can be seen in the formation and breakage of chemical bonds in which energy is released when bonds are formed and energy is absorbed when bonds are broken. At constant temperature and pressure, the equilibrium constant (Keq) is a measure of the directionality of a reaction beginning with equal concentrations of all reactants and products. For example, in a reaction in which the reactants A and B are converted to the products C and D, and a, b, c, and d are the moles of each reactant and product, the Keq is defined by the concentration of each reactant and product, [ ] denotes concentration in moles, when the reaction has reached equilibrium using the relationship:

aA + bB cC + dD Keq = [Ceq]c [Deq]d

[Aeq]a[Beq]b

If the Keq > 0, it means that the reaction proceeds toward product formation (left to right as written), whereas, a Keq < 0 means the reaction favors the formation of reactant (right to left as written). The relationship between the change in free energy, and the reactant and product concentrations under initial conditions, was defined by Gibbs using the following equation in which ΔGº refers to the standard free energy change (kJ/mol), R is the gas constant (8.315 J/mol•K), and T is the absolute temperature:

ΔG = ΔGº + RT ln [Ci] [Di] [Ai] [Bi]


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The concentrations of reactants and products in this equation are given as the initial concentrations ([Ai], [Bi], etc.), which in biological systems, refers to the steady state concentration in living cells (homeostasis). As biochemists, we redefine both the actual ΔG, and the standard change in free energy, ΔGº, to ΔG' and ΔGº', respectively, which refers to free energy changes under physiological conditions (pH 7 and the concentration of H2O as 55.5M).

The importance of the free energy values in biochemistry is that it can be used to predict if a reaction is favorable or unfavorable given the characteristic ΔGº' value for the reaction, and the concentrations of each reactant and product. The ΔGº' value is determined by setting up the reaction under standard physiological conditions (pH7, 55.5M H2O, 298ºK, 1 atm., 1 M of solute), and then allowing it to proceed to equilibrium, at which time, the concentrations of all reactants and products are measured ([Aeq], [Beq], etc.). Once the reaction has reached equilibrium, the change in free energy is zero (ΔG' = 0), and ΔGº' can be calculated directly from the Keq:

0 = ΔGº' + RT ln [Ceq] [Deq] [Aeq][Beq]

ΔGº' = -RT ln Keq

The values of ΔGº' and Keq can both be used to describe the spontaneity of a reaction, in that a favorable reaction with a Keq > 0 corresponds to a ΔGº' < 0, which denotes an exergonic reaction. Similarly, an unfavorable reaction with a Keq < 0 will have a ΔG' > 0 and is endergonic. There are two important points to make about ΔG' and ΔGº' values. First, based on Gibbs equation, the actual free energy change of a reaction inside a living cell can be favorable (ΔG' <0) even if the characteristic ΔGº' for the reaction is unfavorable (ΔGº >0) under equilibrium conditions. This is because metabolic flux in living cells (the flow of metabolites through metabolic pathways) can maintain the steady state concentrations of reactants and products far from their equilibrium concentrations. This is often accomplished using coupled reactions that take advantage of the high free energy content (potential energy) of ATP. Moreover, coupled reactions quickly remove products since these same metabolites serve as reactants in coupled reactions. Importantly, the natural log (ln) of a number that is less than 1 is a negative number. Therefore, even if the standard free energy change is unfavorable, (ΔGº' >0), the overall ΔG' can still be favorable if the mass action ratio, i.e., products]/[reactants], is less than 1 since the ln of this number will be negative. If the concentration of products is very small compared to the concentration of reactants, then the ln • mass action ratio value will be very negative.

ΔG' = ΔGº' + RT ln [products]actual [reactants]actual

Exergonic and endergonic reactions are coupled in metabolism This brings us to the question of how unfavorable (endergonic) reactions in living systems are able to occur. The answer is that endergonic reactions are coupled to exergonic reactions such that the overall change in free energy is favorable (exergonic). The thermodynamic basis for coupling endergonic and exergonic reactions is that the ΔG' for two reactions that share a common intermediate (the product of the first reaction is a reactant in the second reaction), is equal to the sum of the ΔG' values for the two separate reactions as shown below:


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Figure 8.

Glutamine synthetase complex

Enzyme 1 Reaction 1 A B ΔG'1 = +4 kJ/mol (endergonic) Enzyme 2 Reaction 2 B C ΔG'2 = -10 kJ/mol (exergonic) Net reaction A C ΔG'1 + ΔG'2 = ΔG'3 = -6 kJ/mol (exergonic) Although the first reaction is unfavorable (ΔG'>0), the formation of product C occurs because the combined free energies of reactions 1 and 2 are favorable (ΔG'<0). One way to think about how reaction 2 affects reaction 1 in this example is to realize that by quickly converting B into C, the concentration of B is not able to reach its normal equilibrium concentration. This causes the equilibrium of reaction 1 to be shifted toward more product formation in order to replace the B that is continually being consumed by reaction 2. This coupling is the driving force of metabolic flux.

One of the most common types of coupled reactions is one in which the phosphoanhydride bond energy in ATP is used to drive an unfavorable reaction. ATP contains two phosphoanhydride bonds, sometimes referred to as "high energy phosphate bonds" (~P), that can be used as a source of free energy (figure 8). Note however, that the term high energy phosphate bond can be misleading since phosphoanhydride bonds are no different than any other chemical bond and adhere to the laws of thermodynamics. The change in standard free energy (ΔGº') for cleavage of the phosphoanhydride bond between the β and γ phosphate of ATP is -30.5 kJ/mol, and that of the phosphoanhydride bond between the α and β phosphates is -32.3 kJ/mol. Importantly, while hydrolysis reactions were used to calculate these ΔGº' values, it is not the cleavage of these phosphoanhydride bonds that provides energy for coupled metabolic reactions (bond cleavage actually requires energy), but rather the transfer of a phosphate or adenylate (AMP) group to a reactant to generate a highly reactive intermediate. ATP-coupled reactions take place within the active site of an enzyme which accelerates the rate of a reaction by providing an ideal chemical environment for product formation. In these enzyme-mediated coupled reactions, the phosphorylated or adenylated chemical intermediates are the compounds that function as the shared intermediate in the two reactions.

The conversion of the amino acid glutamate to glutamine by the enzyme glutamine synthetase is a


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Fatty acyl-CoA synthetase with ATP

good example of an ATP-coupled reaction that can be broken down into two reactions. The first reaction is endergonic (ΔGº' = +14.2 kJ/mol) and the second reaction is exergonic (ΔGº' = -30.5 kJ/mol) to give an overall favorable ΔGº of -16.3 kJ/mol:

Glutamate + NH4

+ glutamine ΔGº' = +14.2 kJ/mol ATP ADP + Pi ΔGº' = -30.5 kJ/mol Glutamate + NH4

+ + ATP glutamine + ADP + Pi ΔGº' = -16.3 kJ/mol

A second type of ATP-coupled reaction is exemplified by the enzyme fatty acyl-CoA synthetase which catalyzes a coupled reaction that generates an enzyme-linked acyl-adenylate intermediate and the release of pyrophosphate (PPi). The adenylyl group is then replaced with coenzyme A to generate the products fatty acyl-CoA and AMP. An added feature of coupled adenylation transfer reactions is that PPi is rapidly hydrolyzed to 2 Pi by the enzyme inorganic pyrophosphatase which drops the overall ΔGº' for the reaction by another -19.2 kJ/mol. This is important because the formation of a fatty acyl-CoA, such as palmitoyl-CoA, is a highly endergonic reaction with a ΔGº' of +31.4 kJ/mol. In this case, both phosphoanhydride bonds of ATP are ultimately required to drive this reaction to product formation:

Palmitic acid + CoA-SH palmitoyl-CoA ΔGº' = +31.4 kJ/mol ATP AMP + PPi ΔGº' = -30.5 kJ/mol PPi 2 Pi ΔGº' = -19.2 kJ/mol Palmitic acid + ATP palmitoyl-CoA + AMP + 2 Pi ΔGº' = -18.3 kJ/mol

Steady-state substrate concentrations are also a contributing factor to metabolic flux as they reflect the mass action ratio which determines the actual ΔG of a reaction. Remember that the overall ΔG for a metabolic pathway is indeed negative inside the cell as predicted by thermodynamic principles, i.e., life is favorable as long as energy is available to prevent the organism from reaching equilibrium with the environment. With that in mind, let’s see how the actual free energy change ΔG of a reaction is affected by the steady-state concentrations of substrates and products using coupled reactions in the glycolytic pathway. The ΔGº’ of the phosphoglucoisomerase reaction is positive (ΔGº’ = +1.7 kJ/mol). However, since the product of the reaction, fructose-6-P, is quickly converted to fructose-1,6-bisphosphate by the next enzyme in the pathway, phosphofructokinase-1, which is an ATP-coupled reaction (ΔGº’ = -14.2 kJ/mol), then the concentration of fructose-6-P is kept low and the phosphoglucoisomerase reaction is “pulled” to the right under actual conditions (ΔG' = -2.9 kJ/mol). This can be shown by calculating ΔG' using the steady-state concentrations of glucose-6-P and fructose-6-P in erythrocytes at 37ºC as shown below:

Glucose-6-P fructose-6-P ΔGº’ = +1.7 kJ/mol

[glucose-6-P]actual = 8.3 x 10-5 M [fructose-6-P]actual = 1.4 x 10-5M


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Figure 9.

ΔG' = ΔGº’+ RT • ln (mass action ratio) ΔG' = +1.7 kJ/mol + RT • ln [fructose-6-P]actual

[glucose-6-P]actual

ΔG' = +1.7 kJ/mol + (0.00831 kJ/mol • K) (310 ºK) • ln(1.4 x 10-5M) (8.3 x 10-5 M) ΔG' = +1.7 kJ/mol + -4.6 kJ/mol

ΔG' = -2.9 kJ/mol for the reaction Glucose-6-P fructose-6-P

As seen in figure 9, the availability of ATP for the phosphofructokinase-1 reaction turns an unfavorable set of reactions (ΔGº'>0) into a favorable "mini-pathway" by altering the mass action ratio (ΔG'<0). In the absence of ATP, the ΔGº' values for mini-pathway shows that if the reactions are allowed to go to equilibrium, then glucose-6P formation is favored. However, by adding ATP, which contributes phosphoryl transfer energy to the phosphofructokinase-1 reaction, then the steady-state level of fructose-6P is decreased and fructose-1,6-BP is favored.

Note that it is not always possible to measure steady-state levels of metabolites in cells, especially if they are inside organelles such as mitochondria or chloroplast (glycolysis takes place in the cytosol). Therefore, the ΔG' for a reaction cannot be determined and one must assume that the overall reactions in a pathway are favorable even if ΔGº' is not favorable. Nevertheless, often by simply adding up the ΔGº' values for a set of linked reactions, the total ΔGº' is indeed favorable, and even if it isn't, you know that it must be in nature for the reactions to proceed. The adenylate system and the Energy Charge of the cell Since ATP plays such an important role in the cell as a source of free energy for coupled reactions and mechanical work, its levels need to be maintained within a fairly narrow range to avoid a metabolic catastrophe. This is done by interconverting ATP, ADP, and AMP using several key phosphoryl transfer reactions, that together, constitute the adenylate system. To see why the adenylate system is important, consider that a 70 kg person requires ~100 moles of ATP every day based on the energy content of food. The molecular weight of ATP is 507g/mol, which means we hydrolyze as much as 50 kg of ATP every day. Rather than synthesizing our own weight in ATP on a daily basis, it is much more efficient to recycle adenylate forms by reforming ATP from ADP, AMP, and Pi. The most common way is for ADP + Pi to be converted to ATP by the enzyme ATP synthase which is a component of the aerobic respiration and photosynthesis pathways. Both aerobic respiration and photosynthesis require an input of energy in order to drive the ATP synthase reaction forward, i.e., metabolic fuel and light, respectively. Note that since some


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Figure 10.

reactions lead to production of AMP rather than ADP, for example, the adenylation transfer reaction catalyzed by fatty acyl-CoA synthetase, AMP must first be phosphorylated to generate ADP. This reaction is catalyzed by the enzyme adenylate kinase which transfers the γ-phosphate from ATP to AMP to generate 2 ADP (ΔGº' = ~0 kJ/mol). The 2 ADP generated by adenylate kinase can then be phosphorylated to 2 ATP by ATP synthase (ΔGº' for 2 ATP = 61 kJ/mol) to give the net reaction of AMP + 2 Pi --> ATP as shown below: (Adenylate kinase) AMP + ATP 2 ADP (ATP synthase) 2 ADP + 2 Pi ATP + ATP

Net: AMP + 2 Pi ATP

Since ATP is the high energy form of the adenylate system, then the ratio of the concentration of ATP to the concentration of ADP and AMP in the cell at any given time can be used as a measure of the energy state of the cell. This relationship can be expressed in terms of the Energy Charge (EC) which takes into account the number of phosphoanhydride bonds available for work:

Energy Charge (EC) = [ATP] + 0.5 [ADP] [ATP] + [ADP] + [AMP]

If the adenylate system components were present in the cell at the same concentration such that [ATP] = [ADP] = [AMP], then EC = 0.5. However, most cells are found to have an EC in the range of 0.7 to 0.9 which means that the [ATP] is higher than [ADP] or [AMP] (figure 10). For example, under steady state conditions, the EC value in rat hepatocytes is 0.8 based on adenine nucleotide concentrations of:

[ATP] = 3.4 mM [ADP] = 1.3 mM [AMP] = 0.3 mM

Energy Charge (EC) = [ATP] + 0.5 [ADP] [ATP] + [ADP] + [AMP] Energy Charge (EC) = 3.4 mM + 0.5 (1.3 mM) = 0.8 3.4 mM + 1.3 mM + 0.3 mM

The EC of a cell is maintained between 0.7 and 0.9 by regulating metabolic flux through

pathways that generate and consume ATP. Photosynthetic autotrophs use sunlight as their source of energy for ATP production, whereas heterotrophs, use nutrients present in their diet as a source of metabolic fuel in the form of carbohydrate, protein, and lipid, to synthesize ATP. Most organisms use stored metabolic fuels as a source of energy when other forms of energy are not readily available. Extracting energy from metabolic fuels is the function of catabolic pathways


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Figure 11.

Figure 12.

which convert energy-rich compounds into energy-depleted compounds, and in the process, generate reduced coenzymes (NADH, NADPH, and FADH2), as well as, ATP (figure 11). These reduced coenzymes and ATP are then used for the biosynthesis of biomolecules through anabolic pathways. In general, catabolic pathways are degradative processes that obtain energy from compounds using redox reactions to generate ATP, while anabolic pathways are biosynthetic processes that utilize energy from redox reactions and ATP to restrain entropy and maintain order.

Importantly, regulatory processes control the activity of key enzymes in catabolic and anabolic metabolic pathways in order to stabilize the EC and maintain homeostasis. For example, when EC levels decrease due to increased rates of aerobic respiration or sustained flux through anabolic pathways, then enzymes responsible for ATP synthesis become activated and flux through catabolic pathways is increased (figure 12). In most organisms, this means degrading stored metabolic fuel in the form of carbohydrate or lipid. Alternatively, when EC levels are elevated due to photosynthesis or high levels of nutrients following a meal, then enzymes that control flux through anabolic pathways are activated to take advantage of the available ATP and replenish stored metabolic fuel.


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ANSWERS TO KEY CONCEPT QUESTIONS: How is energy from the sun converted to chemical energy? Life on earth is made possible by the biochemical reactions of photosynthesis, carbon fixation and aerobic respiration which together convert solar energy into ATP (and NADPH) which is used to synthesize carbohydrates from CO2 and H2O. Aerobic organisms, such as ourselves, consume carbohydrates as a chemical source of energy and metabolize them in the presence of O2 to from CO2 and H2O. All organisms depend directly or indirectly on energy derived from thermonuclear fusion reactions on the sun to prevent (for as long as possible) reaching equilibrium with the environment; a high entropy state called death. What is reaction coupling and why is it important in metabolic pathways? Reaction coupling permits energetically unfavorable reactions to be more favorable in the context of a pathway. Coupling ATP hydrolysis to a phosphoryl transfer reaction is one example of reaction coupling that takes place in the same enzyme active site. The net ΔGº’ for an ATP coupled reaction is often highly negative, for example, the phosphorylation of glucose by the enzyme hexokinase. Another type of reaction coupling is when two enzymes in pathway are energetically linked through a shared common intermediate. Since the actual change in free energy, ΔG, is the sum of the change in standard free energy, ΔGº’, and RT•ln(mass action ratio), in which the mass action ratio is [product]actual/[substrate]actual , depletion of a reaction product by its metabolism as a substrate in the coupled reaction, results in a reduction in ΔG since ln of a mass action ratio <1 is a negative number. Therefore, even though the ΔGº’ for a reaction is a positive number (based on the reaction reaching equilibrium in a test tube under ideal conditions), the actual ΔG is a negative number because RT•ln(mass action ratio) is a negative number due to lower than "expected" [product] in the cell due to its function as a substrate in a linked reaction of the pathway.

lecture 21 - introduction to metabolism:...

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