Human Nutrition and Metabolism Introduction PDF

Human Nutrition and Metabolism Introduction Introduction Catabolism is the degradative phase of metabolism in which organic nutrient molecules (carbohydrates, fats, and proteins) are converted into smaller, simpler end products (such as lactic acid, CO2, and NH3). Catabolic pathwa...

Human Nutrition and Metabolism Introduction Introduction Catabolism is the degradative phase of metabolism in which organic nutrient molecules (carbohydrates, fats, and proteins) are converted into smaller, simpler end products (such as lactic acid, CO2, and NH3). Catabolic pathways release energy, some of which is conserved in the formation of ATP and reduced electron carriers (NADH or NADPH); the rest is lost as heat. In anabolism, also called biosynthesis, small, simple precursors are built up into larger and more complex molecules, including lipids, polysaccharides, proteins, and nucleic acids. Anabolic reactions require an input of energy, generally in the form of the phosphoryl group transfer potential of ATP and the reducing power of NADH and NADPH. Introduction Most cells have the enzymes to carry out both the degradation and the synthesis of the important categories of biomolecules — fatty acids, for example. The simultaneous synthesis and degradation of fatty acids would be wasteful, however. This is prevented by reciprocally regulating the anabolic and catabolic reaction sequences: when one sequence is active, the other is suppressed. Such regulation could not occur if anabolic and catabolic pathways were catalyzed by exactly the same set of enzymes, operating in one direction for anabolism, the opposite direction for catabolism: inhibition of an enzyme involved in catabolism would also inhibit the reaction sequence in the anabolic direction. Catabolic and anabolic pathways that connect the same two endpoints (glucose →→ pyruvate, and pyruvate →→ glucose, for example) may employ many of the same enzymes; but at least one of the steps is catalyzed by different enzymes in the catabolic and anabolic directions, and these enzymes are the sites of separate regulation. Moreover, for both anabolic and catabolic pathways to be essentially irreversible, the reactions unique to each direction must include at least one that is thermodynamically very favorable — in other words, a reaction for which the reverse reaction is very unfavorable. Introduction As a further contribution to the separate regulation of catabolic and anabolic reaction sequences, paired catabolic and anabolic pathways commonly take place in different cellular compartments: for example, fatty acid catabolism occurs in animal mitochondria, whereas fatty acid synthesis occurs in the cytosol. The concentrations of intermediates, enzymes, and regulators can be maintained at different levels in these different compartments. Because metabolic pathways are subject to kinetic control by substrate concentration, separate pools of anabolic and catabolic intermediates also contribute to the control of metabolic rates. Introduction Metabolic pathways are regulated at several levels, from within the cell and from outside. A key enzyme in a pathway may be activated allosterically, or its amount may be changed by the rates of synthesis and breakdown of the enzyme. In multicellular organisms, the metabolic activities of different tissues are regulated and integrated by growth factors and hormones that act from outside the cell. Introduction Living cells and organisms must perform work to stay alive, to grow, and to reproduce. The ability to harness energy and to channel it into biological work is a fundamental property of all living organisms. Modern organisms carry out a remarkable variety of energy transductions, conversions of one form of energy to another. They use the chemical energy in fuels to bring about the synthesis of complex, highly ordered macromolecules from simple precursors. They also convert the chemical energy of fuels into concentration gradients and electrical gradients, into motion and heat, and, in a few organisms such as fireflies and some deep-sea fish, into light. Introduction The chemical changes and energy transductions in living organisms follow the laws of thermodynamics. The free-energy change is the maximum energy made available to do work when a chemical reaction occurs. If two reactions can be combined to yield a third reaction, the overall free energy change is the sum of the two. Cells accomplish energy-requiring chemical work by coupling an energy-releasing (exergonic) reaction such as the cleavage of ATP to an endergonic reaction (which requires energy input). ATP is the universal energy currency in living organisms Transfer of its phosphoryl group to a water molecule or metabolic intermediates provides the energetic push for muscle contraction, the pumping of solutes against concentration gradients, and the synthesis of complex molecules. Introduction Oxidation-reduction reactions indirectly provide much of the energy needed to make ATP. Reduced substrates such as glucose are oxidized in several steps, with the energy of oxidation steps conserved in the form of a reduced cofactor, NADH. Energy stored in NADH is used to drive the synthesis of ATP. To respond to changes in external circumstances, cells must regulate enzyme activities, by changing either the number of enzyme molecules or the catalytic activity of preexisting enzyme molecules. Introduction 13.1 Bioenergetics and Thermodynamics Bioenergetics is the quantitative study of energy transductions — changes of one form of energy into another — that occur in living cells, and of the nature and function of the chemical processes underlying these transductions. Introduction Biological Energy Transformations Obey the Laws of Thermodynamics Many quantitative observations made by physicists and chemists on the interconversion of different forms of energy led, in the nineteenth century, to the formulation of two fundamental laws of thermodynamics. The first law is the principle of the conservation of energy: for any physical or chemical change, the total amount of energy in the universe remains constant; energy may change form or it may be transported from one region to another, but it cannot be created or destroyed. The second law of thermodynamics, which can be stated in several forms, says that the universe always tends toward increasing disorder: in all natural processes, the entropy of the universe increases. Introduction The reacting system is the collection of matter that is undergoing a particular chemical or physical process; it may be an organism, a cell, or two reacting compounds. The reacting system and its surroundings together constitute the universe. In the laboratory, some chemical or physical processes can be carried out in isolated or closed systems, in which no material or energy is exchanged with the surroundings. Living cells and organisms, however, are open systems, exchanging both material and energy with their surroundings; living systems are never at equilibrium with their surroundings, and the constant transactions between system and surroundings explain how organisms can create order within themselves while operating within the second law of thermodynamics. Introduction Below are the three thermodynamic quantities that describe the energy changes occurring in a chemical reaction: Free energy, G (for J. Willard Gibbs), expresses the amount of energy capable of doing work during a reaction at constant temperature and pressure. When a reaction proceeds with the release of free energy (that is, when the system changes so as to possess less free energy), the free energy change, ΔG, has a negative value and the reaction is said to be exergonic. In endergonic reactions, the system gains free energy and ΔG is positive. Introduction Enthalpy, H, is the heat content of the reacting system. It reflects the number and kinds of chemical bonds (covalent and noncovalent) in the reactants and products. When a chemical reaction releases heat, it is said to be exothermic; the heat content of the products is less than that of the reactants, and the change in enthalpy, ΔH, has, by convention, a negative value. Reacting systems that take up heat from their surroundings are endothermic and have positive values of ΔH. Introduction Entropy, S, is a quantitative expression for the randomness or disorder in a system. When the products of a reacting system are less complex and more disordered than the reactants, the reaction is said to proceed with a gain in entropy. The units of ΔG and ΔH are joules/mole or calories/mole (recall that 1 cal =4.184 J); units of entropy are joules/mole Kelvin (J/mol ∙ K). Introduction Under the conditions existing in biological systems (including constant temperature and pressure), changes in free energy, enthalpy, and entropy are related to each other quantitatively by the equation Introduction in which ΔG is the change in Gibbs free energy of the reacting system, ΔH is the change in enthalpy of the system, T is the absolute temperature, and ΔS is the change in entropy of the system. By convention, ΔS has a positive sign when entropy increases and ΔH, as noted above, has a negative sign when heat is released by the system to its surroundings. Either of these conditions, both of which are typical of energetically favorable processes, tends to make ΔG negative. In fact, ΔG of a spontaneously reacting system is always negative. Introduction Cells are isothermal systems — they function at essentially constant temperature (and also function at constant pressure). Heat flow is not a source of energy for cells, because heat can do work only as it passes to a zone or an object at a lower temperature. The energy that cells can and must use is free energy, described by the Gibbs free-energy function G, which allows prediction of the direction of chemical reactions, their exact equilibrium position, and the amount of work they can (in theory) perform at constant temperature and pressure. Heterotrophic cells acquire free energy from nutrient molecules, and photosynthetic cells acquire it from absorbed solar radiation. Both kinds of cells transform this free energy into ATP and other energy-rich compounds capable of providing energy for biological work at constant temperature. Introduction In living cells, reactions that would be extremely slow if uncatalyzed are caused to proceed not by supplying additional heat but by lowering the activation energy through use of an enzyme catalyst. An enzyme provides an alternative reaction pathway with a lower activation energy than the uncatalyzed reaction, so that at body temperature a large fraction of the substrate molecules have enough thermal energy to overcome the activation barrier, and the reaction rate increases dramatically. The free-energy change for a reaction is independent of the pathway by which the reaction occurs; it depends only on the nature and concentration of the initial reactants and the final products. Enzymes cannot, therefore, change equilibrium constants; but they can and do increase the rate at which a reaction proceeds in the direction dictated by thermodynamics (see Section 6.2). Introduction Standard Free-Energy Changes Are Additive In the case of two sequential chemical reactions, A ⇌ B and B ⇌ C, each reaction has its own equilibrium constant and each has its characteristic standard free-energy change, ΔG′1° and ΔG′2°. As the two reactions are sequential, B cancels out to give the overall reaction A ⇌ C, which has its own equilibrium constant and thus its own standard free-energy change, ΔG′ ‹ Sum. Introduction Introduction Introduction This principle of bioenergetics explains how a thermodynamically unfavorable (endergonic) reaction can be driven in the forward direction by coupling it to a highly exergonic reaction. For example, in many organisms, the synthesis of glucose 6-phosphate is the first step in the utilization of glucose. In principle, the synthesis could be accomplished by this reaction: But the positive value of ΔG′° predicts that under standard conditions the reaction will tend not to proceed spontaneously in the direction written. Another cellular reaction, the hydrolysis of ATP to ADP and Pi, is highly exergonic: Introduction But the positive value of ΔG′° predicts that under standard conditions the reaction will tend not to proceed spontaneously in the direction written. Another cellular reaction, the hydrolysis of ATP to ADP and Pi, is highly exergonic: Introduction These two reactions share the common intermediates Pi and H2O and may be expressed as sequential reactions: Introduction The overall standard free-energy change is obtained by adding the ΔG′° values for individual reactions: Introduction The overall reaction is exergonic. In this case, energy stored in ATP is used to drive the synthesis of glucose 6- phosphate, even though its formation from glucose and inorganic phosphate (Pi) is endergonic. The pathway of glucose 6- phosphate formation from glucose by phosphoryl transfer from ATP is different from reactions (1) and (2), but the net result is the same as the sum of the two reactions. The standard free-energy change is a state function. In thermodynamic calculations, all that matters is the state of the system at the beginning of the process and its state at the end; the route between the initial and final states is immaterial. Introduction SUMMARY 13.1 Bioenergetics and Thermodynamics Bioenergetics is the quantitative study of energy relationships and energy conversions in biological systems. Biological energy transformations obey the laws of thermodynamics. Living cells constantly perform work. They require energy for maintaining their highly organized structures, synthesizing cellular components, transporting small molecules and ions across membranes, and generating electric currents. All chemical reactions are influenced by two forces: the tendency to achieve the most stable bonding state (for which enthalpy, H, is a useful expression) and the tendency to achieve the highest degree of randomness, expressed as entropy, S. The driving force in a reaction is ΔG, the free-energy change, which represents the net effect of these two factors: ΔG = ΔH − T ΔS. The standard transformed free-energy change, ΔG′°, is a physical constant that is characteristic for a given reaction and can be calculated from the equilibrium constant for the reaction: ΔG′° = −RT ln K′ eq. Introduction The actual free-energy change, ΔG, is a variable that depends on ΔG′° and on the concentrations of reactants and products: ΔG = ΔG′° + RT ln([products]/[reactants]). When ΔG is large and negative, the reaction tends to go in the forward direction; when ΔG is large and positive, the reaction tends to go in the reverse direction; and when ΔG = 0, the system is at equilibrium. The free-energy change for a reaction is independent of the pathway by which the reaction occurs. Free-energy changes are additive; the net chemical reaction that results from successive reactions sharing a common intermediate has an overall free energy change that is the sum of the ΔG values for the individual reactions. Oxidation-reduction reactions involve the loss or gain of electrons: one reactant gains electrons and is reduced, while the other loses electrons and is oxidized. Oxidation reactions generally release energy and are important in catabolism.