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A Novel Theory for ATP Synthesis
By Baltazar D. Reynafarje
Trafford PublishingCopyright © 2014 Baltazar D. Reynafarje
All rights reserved.
ENERGY TRANSFORMATIONS IN LIVING ORGANISMS
I. 1. Introduction
The central theme of bioenergetics has to do with the principles that govern the transformations of energy in living organisms. Energy, i.e. the ability of matter to perform work as the results of motion or position, exists in various forms, including mechanical, thermal, chemical, electrical, radiant, and atomic. All forms of energy, except entropy, may be converted to other forms of energy. Kinetic energy (the energy related with motion) and potential energy (the energy related to position) may be lost or gained, but the total of the two remains constant.
The concept of equivalence between mass and energy refers to the physical principle establishing that a measured quantity of energy (E) is equivalent to a measured quantity of mass is expressed by the following Einstein's equation:
E = mc2 = (ma).d (1)
In this form of kinetic energy, m is the unified body of mass that occupies space, a is the acceleration applied to m, d is the distance through which a acts, and c is the speed of light. This concept, however, only holds true for events involving velocities equal to the velocity of light. For velocities lower than the speed of light the value of E can be derived from the following equation:
E = ½ mv2 (2)
where v2 is the speed multiplied by itself. At higher velocities close to that of light energy and matter are interconvertible.
Living organisms require a continual input and utilization of free energy for three main purposes: the performance of dynamical energy for cellular motion; the active transport of molecules and ions; and the synthesis of macromolecules from simple precursors. The free-energy that sustains life on both phototropic and chemotropic organisms, comes from solar energy. In plants and photosynthetic bacteria, the protein complexes in charge to make available the quanta of visible light are contained in chloroplasts. In aerobic organisms, the protein complexes in charge to make available the free energy of electron flow towards oxygen are contained in the mitochondria of eukaryotic organisms and the membranes of cyanobacteria and prokaryotic cells
Although the number of protein-assemblies that exists inside the cell is overwhelmingly large, the number of assemblies involved in the synthesis of ATP, which is the universal form of free energy in living organisms, is relatively small. The essential elements that, together with O2, CO2, ADP and Pi, are involved in the oxidative phosphorylation process of ATP synthesis are the electrons and protons contained in highly reduced respiratory substrates. The protein complexes in charge to catalyze the downhill or exergonic flow of electrons towards O2 and the uphill or endergonic synthesis of ATP from ADP and Pi are localized in the inner mitochondrial membrane of eukaryotes.
I. 2. Basic laws of thermodynamics
All biological reactions that occur in nature are subject to the universal laws of thermodynamics. Although cells are open systems that are never at equilibrium, a description of the principles of equilibrium thermodynamics (as applied to ordinary chemical reactions) is necessary to understand that the work done by the cell always takes place under conditions that are far from equilibrium. Thermodynamics is the field of physics that deals with the relationship between heat and other forms of energy such as pressure, temperature and volume. Everything that happens in living organisms is subject to the laws of thermodynamics.
The first law of thermodynamics is the law of energy conservation. It states that energy can neither be created nor destroyed. In any given process, one form of energy may be converted into another but the total energy of the system (the cell) plus its environment remains constant. The total energy of the universe is constant. Whatever its nature (thermal, chemical, electrical, mechanical, kinetic, potential, or atomic), the energy is neither created nor destroyed; it is only transformed and distributed between the system and its universe. A system in a given state has a definite amount of internal energy, i.e. the total kinetic and potential energy associated with the motions and relative positions of the molecules of an object, excluding the kinetic or potential energy of the object as a whole (the system and its environment). A rise in temperature or change in phase of the system, results in a rise of internal energy, generally represented by U. The internal energy can be changed in only two ways: (1) heat energy can flow into or out of the system, and (2) the system can do work of some kind against external forces. Thus,
ΔU = ΔQ - ΔW (3)
where ΔU is the general representation of the change in internal energy of the system, ΔQ is the heat that flow into or out of the system, and ΔW is the work (quantity equal to the force applied to an object time the motion of the object in direction of the applied force) done by the system. By convention, ΔW appears with a negative sign because any work done by the system reduces its internal energy. In reactions in which there is a change in the electrical or oxidation-reduction potential of the system the change in internal energy (ΔE) is mathematically expressed by the following equation:
ΔE (EB – EA)=ΔQ - ΔW (4)
in which EA is the energy of the system at the start of the process and EB the energy of the system at the end of the process. The sign of ΔE depends on the extent of heat that flows into or out of the system. When the heat that flows into the system (ΔQ) increases the internal energy increases and the sign of ΔE in equation 4 is positive.
The first law of thermodynamics is simply a law of conservation of energy. Nothing is said about the relative usefulness, direction or spontaneity of the reaction. Some reactions do occur spontaneously even when ΔE is positive, i.e. when the free energy of the system increases. In such cases, the system absorbs heat from its surroundings and the entropy (S) or degree of randomness or disorder of the system decreases.
The second law of thermodynamics states that all naturally occurring processes proceed toward equilibrium, i.e. in the direction of minimal potential energy. These so-called spontaneous reactions release energy that can be harnessed, transformed and made to do work. A more complete statement of the second law of thermodynamics includes the concept of entropy (S), which in a closed thermodynamic system is a quantitative measure of the amount of thermal energy not available to do work. Thus, in accordance with the second law of thermodynamics a reaction can occur spontaneously only if the sum of the entropies of the system and its surrounding increase or is higher than zero:
(ΔSsystem + ΔSsurroundings) > 0 (5)
The more random, disordered, disorganized, or chaotic the system, the higher is the entropy. Spontaneous processes only occur when the entropy of the system and its environment increases. For example, the hydrolysis of ATP to Pi and ADP is thermodynamically feasible because the entropy of the system, ATP, is smaller than the entropy of Pi, ADP, and the surroundings. The energy lost by the hydrolysis of ATP (-7.3 kcal/mol) is lower than the actual energy required for its synthesis (ΔG = -12 kcal/mol). While the total energy of a system and its surrounding remains constant, the energy is distributed in a quantitatively different way after a spontaneous reaction.
The second law of thermodynamics, however, says nothing about the extent and relative utility of the transformations of energy. If matter and energy can only go from a state of maximal organization to a state of inert uniformity, maximal disorganization and chaos, how then can we explain the very existence of life, the quality that distinguishes highly organized matter from inanimate and disorganized matter? Life, however, does not violate any law of thermodynamics. The natural tendency of matter and energy in a given organism to run downhill can be counteracted by putting energy or doing work on the organism. In any particular system the energy may increase, remain constant or decrease while the total energy of the system, i.e. the energy of the system receiving the energy plus that of the system providing the energy, remains constant. A more complete statement of the second law that takes into account the single-directionality of spontaneous processes and the decreased potential to do further work is this: "the entropy of the universe is constantly increasing".
The third law of thermodynamics states that at a temperature of absolute zero (0°K), where all random motion ceases, the entropy of a perfect crystal is zero, that is, all the atoms are maximally organized. However, if in accordance with the laws of thermodynamics the randomness or disorder of the universe is constantly increasing, do we have to assume that the universe was at some time a perfect crystal? Consider the adequacy of the Big Bang theory.
I. 3. Endergonic, exergonic, and Gibbs free energy change in living organisms
Living cells are exceedingly complex and delicate structures that grow and multiply maintaining their integrity over long periods of time by utilizing the energy contained in "energy-rich compounds". Chemical reactions that yield free energy and are capable of doing work are called "exergonic". Those that utilize energy and need work to be done in order to proceed are called "endergonic". Seldom, however, the free energy released by an exergonic reaction is directly transferred to an endergonic reaction. In general, the free energy released during exergonic reactions is conserved through a series of coupled reactions until the final endergonic reaction takes place. For example, the free energy released during the exergonic oxidation of respiratory substrates, is first transformed into the free energy of electron flow which eventually is utilized in the endergonic process of ATP synthesis.
Even though the entropy of the universe is constantly increasing, exergonic and endergonic reactions in closed thermodynamic systems can apparently happen "spontaneously". One difficulty in using entropy to define a spontaneous reaction is that the entropy is practically impossible to measure (see equation 4). This difficulty is obviated by using the concept of free energy change or Gibbs free energy change (ΔG), as represented by the following equation:
ΔG = ΔH - TΔS (6)
In this equation, ΔG is the change in free energy of a system undergoing transformation at constant pressure (P) and temperature (T), ΔH is the change in enthalpy or heat content of the system, and ΔS is the change in entropy,
ΔH = ΔE + PΔV (7)
Because the volume-change (ΔV) in almost every biochemical reaction is very small and ΔH is nearly equal to ΔE, the change in free energy of the system (ΔG) can be approximately represented by the following equation:
ΔG = ΔE - TΔS (8)
A reaction can occur spontaneously only when the change in entropy or disorder in the system (TΔS,) increases above the energy content of the system (ΔE) and the ΔG is negative. If the ΔG is positive the reaction will not occur spontaneously unless an input of free energy from a coupled reaction drives the system towards equilibrium. When a physical system moves from one state of equilibrium to another, a thermodynamic process is said to take place. This phenomenon is divided into the part being studied, the system, and the region around the system, the surroundings. Whenever work is done on an object, there is a transfer of energy to the object, and so work is considered to be energy in transit. If a constantly acting force does not produce motion there is no work performed. For example, steadily holding an object above the floor does not involve any work, but the energy used to lift the object from the floor is retained in the object. This energy is released only when the object falls to the floor. The energy released or utilized in a chemical reaction at constant temperature and pressure is called free energy difference, ΔG, or Gibbs free energy change. The G of a chemical reaction represents the difference between the ratio of products and substrates at the beginning of the reaction and the ratio of products and substrates at equilibrium.
ΔG = [actual product/substrate ratios] - [product/substrate ratios at equilibrium] (9)
More specifically, in the reaction from A and B (substrates) to C and D (products)
ΔG = actual RT ln [C]c [D]d / [A]a [B]b - RT ln [C]c [D]d / [A]a [B]b at K'eq (10)
Where R is the gas constant = 1.987 cal mole-1 °K-1, T is the absolute temperature, °K at 25°C = 298°K, and a, b, c, d = coefficients of A, B, C and D, respectively.
Excerpted from Bioenergetics by Baltazar D. Reynafarje. Copyright © 2014 Baltazar D. Reynafarje. Excerpted by permission of Trafford Publishing.
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Table of Contents
ContentsSECTION I: ENERGY TRANSFORMATIONS IN LIVING ORGANISMS, 1,
I. 1. Introduction, 1,
I. 2. Basic laws of thermodynamics, 3,
I. 3. Endergonic, exergonic, and Gibbs free energy change in living organisms, 6,
I. 4. Relationship between free-energy change and equilibrium constant, 10,
I. 5. Enthalpy and Entropy, 13,
SECTION II: THERMODYNAMICALLY FAVORABLE AND UNFAVORABLE REACTIONS, 14,
II. 1. Introduction, 14,
II. 2. The active transport of ions and metabolites requires the input of free energy, 15,
II. 3. Most important forms of metabolic energy, 22,
II. 4. Electrons enter the mitochondria via specific shuttles, 28,
II. 5. Acetyl CoA is the common entrance of electrons to the Krebs cycle, 36,
SECTION III: KINETICS OF ENZYME CATALYZED REACTIONS, 42,
III. 1. Introduction, 42,
III. 2. Kinetic orders and steady-state conditions, 43,
III. 3. The Michaelis-Menten constant of hyperbolical reactions, 47,
III. 4. Turnover number and Hill coefficient of enzyme catalyzed reactions, 52,
SECTION IV: TRANSPORT OF OXYGEN FROM AIR TO TISSUES, 60,
IV. 1. Introduction, 60,
IV. 2. Transport of oxygen by the cardiovascular system, 62,
IV. 3. Physiological transport of O2 form blood to mitochondria, 65,
SECTION V: THE RESPIRATORY ROCESS OF O2 CONSUMPTION, 70,
V. 1. Introduction, 70,
V. 2. Components of the mitochondrial membranes, 70,
V. 3. Complexes of the respiratory chain, 73,
V. 4. 1.14 volts drives the transfer of electrons from NADH to O2 via QH2, 75,
V. 5. Cytochrome aa3 catalyzes the transfer of electrons from cytochrome c to O2,
V. 6. Metabolic states of mitochondria and the concept of respiratory control, 82,
SECTION VI: KINETICS AND THERMODYNAMICS OF OXYGEN CONSUMPTION, 87,
VI. 1. Introduction, 87,
VI. 2. The KM of Cytochrome Oxidase for Oxygen is close to 30 µM, 88,
VI. 3. Effect of the O/cytochrome aa3 ratio on the extent of O2 consumption, 90,
VI. 4. The vectorial ejection of H+ has no effect on the process of ATP synthesis, 93,
SECTION VII: KINETICS AND THERMODYNAMICS OF ATP SYNTHESIS, 96,
VII. 1. Introduction, 96,
VII. 2. The synthesis of ATP depends more on O2 than on ADP concentration, 97,
VII. 3. The efficiency of the mitochondrial process of ATP synthesis, 100,
VII. 4. Net synthesis of ATP takes place in the absence of a protonmotive force, 102,
VII. 5. Distinct mechanisms for the synthesis and hydrolysis of ATP, 104,