Basic Thermodynamics PDF

Basic Thermodynamics  Physical Foundations in Biological Chemistry One of the goals is to understand, in quantitative and chemical terms, the means by which energy is extracted, channeled, and consumed in living cells. cellular energy conversions—like all other ener...

Basic Thermodynamics  Physical Foundations in Biological Chemistry One of the goals is to understand, in quantitative and chemical terms, the means by which energy is extracted, channeled, and consumed in living cells. cellular energy conversions—like all other energy conversions—can be considered in the context of the laws of thermodynamics. Small molecules, macromolecules, and supramolecular complexes are continuously synthesized and broken down in chemical reactions that involve a constant flux of mass and energy through the system. The amounts of hemoglobin and glucose in the blood remain nearly constant because the rate of synthesis or intake of each just balances the rate of its breakdown, consumption, or conversion into some other product. The constancy of concentration is the result of a dynamic steady state, a steady state that is far from equilibrium. Maintaining this steady state requires the constant investment of energy; when a cell can no longer generate energy, it dies and begins to decay toward equilibrium with its surroundings. How is the energy interconversions in living organisms occurs in the system and surroundings? Isolated System If the system exchanges neither matter nor energy In a chemical reactions occurring in with its surroundings solution, all the constituent reactants and products, the solvent that contains them, Closed and the immediate Atmosphere —in short, everything within a defined region of space If the system exchanges Universe = + energy but not matter with its surroundings The system and its Open surroundings together Surroundings constitute the universe if the system exchanges Anything other than both energy and matter system with its surroundings A living organism is an open system; it exchanges both matter and energy with its surroundings. Organisms derive energy from their surroundings in two ways: (1) they take up chemical fuels (such as glucose) from the environment and extract energy by oxidizing them. (2) they absorb energy from sunlight. The first law of thermodynamics describes the principle of the conservation of energy: in any physical or chemical change, the total amount of energy in the universe remains constant, although the form of the energy may change Energy interconversions in living organisms  Entropy The key descriptors of entropy are randomness and disorder, manifested in different ways. 1. The Teakettle and the Randomization of Heat we turn off the burner under a tea kettle full of water at 100 C (the “system”) in the kitchen (the “surroundings”) and allow the teakettle to cool. As it cools, no work is done, but heat passes from the teakettle to the surroundings, raising the temperature of the surroundings (the kitchen) by an infinitesimally small amount until complete equilibrium is attained. The free energy that was once concentrated in the teakettle of hot water at 100 C, potentially capable of doing work, has disappeared. Its equivalent in heat energy is still present in the teakettle kitchen (i.e., the “universe”) but has become completely randomized throughout. This energy is no longer available to do work because there is no temperature differential within the kitchen. Moreover, the increase in entropy of the kitchen (the surroundings) is irreversible 2: The Oxidation of Glucose C6H12O6 + 6O2 6CO2 + 6H2O The end products of this oxidative metabolism, CO2 and H2O, are returned to the surroundings. In this process the surroundings undergo an increase in entropy, whereas the organism itself remains in a steady state and undergoes no change in its internal order Whenever a chemical reaction results in an increase in the number of molecules—or when a solid substance is converted into liquid or gaseous products, which allow more freedom of molecular movement than solids— molecular disorder, and thus entropy, increases  The Flow of Electrons Provides Energy for Organisms Nearly all living organisms derive their energy, directly or indirectly, from the radiant energy of sunlight light-driven splitting of water during photosynthesis releases its electrons for the reduction of CO2 and the release of O2 into the atmosphere: Nonphotosynthetic cells and organisms obtain the energy they need by oxidizing the energy-rich products of photosynthesis, then passing the electrons thus acquired to atmospheric O2 to form water, CO2, and other end products, which are recycled in the environment Thus autotrophs and heterotrophs participate in global cycles of O2 and CO2, driven ultimately by sunlight, making these two large groups of organisms interdependent. All these reactions involved in electron flow are oxidation reduction reactions: one reactant is oxidized (loses electrons) as another is reduced (gains electrons) Creating and Maintaining Order Requires Work and Energy According to the second law of thermodynamics, the tendency in nature is toward ever-greater disorder in the universe: the total entropy of the universe is continually increasing. To bring about the synthesis of macromolecules from their monomeric units, free energy must be supplied to the system (in this case, the cell). The randomness or disorder of the components of a chemical system is expressed as entropy, S. Any change in randomness of the system is expressed as entropy change, ΔS, which by convention has a positive value when randomness increases. Free-energy content, G, of any closed system can be defined in terms of three quantities: enthalpy, H, reflecting the number and kinds of bonds; entropy, S; and the absolute temperature, T (in Kelvin). These variables are related by the equation ΔG = ΔH - T ΔS. When reaction is occuring at constant temp. i.e. T = 1, then ΔG, is determined by the enthalpy change, ΔH, reflecting the kinds and numbers of chemical bonds and noncovalent interactions broken and formed, and the entropy change, ΔS, describing the change in the system’s randomness. ΔH is negative for a reaction that releases heat, and ΔS is positive for a reaction that increases the system’s randomness A process tends to occur spontaneously only if ΔG is negative (if free energy is released in the process). To carry out these thermodynamically unfavorable, energy-requiring (endergonic) reactions, cells couple them to other reactions that liberate free energy (exergonic reactions), so that the overall process is exergonic : the sum of the free-energy changes (ΔG1 + ΔG2)is negative. By this coupling strategy, cells are able to synthesize and maintain the information-rich polymers essential to life. Energy coupling in mechanical and chemical processes  K eq and ΔG0 Are Measures of a Reaction’s Tendency to Proceed Spontaneously The tendency of a chemical reaction to go to completion can be expressed as an equilibrium constant. For the reaction in which a moles of A react with b moles of B to give c moles of C and d moles of D, the equilibrium constant, Keq, is given by where [A]eq is the concentration of A, [B]eq the concentration of B, and so on, when the system has reached equilibrium. A large value of Keq means the reaction tends to proceed until the reactants are almost completely converted into the products. Gibbs showed that ΔG (the actual free-energy change) for any chemical reaction is a function of the standard free- energy change, ΔG0 —a constant that is characteristic of each specific reaction—and a term that expresses the initial concentrations of reactants and products: where [A]i is the initial concentration of A, and so forth; R is the gas constant; and T is the absolute temperature. ΔG is a measure of the distance of a system from its equilibrium position. When a reaction has reached equilibrium, no driving force remains and it can do no work: ΔG = 0. For this special case, [A]i = [A]eq, and so on, for all reactants and products, and Substituting 0 for ΔG and Keq for in Equation 1, we obtain the relationship from which we see that ΔG0 is simply a second way (besides Keq) of expressing the driving force on a reaction. Because Keq is experimentally measurable, we have a way of determining ΔG0 , the thermodynamic constant characteristic of each reaction. The units of ΔG and ΔG0 are joules per mole (or calories per mole). When Keq>>>1, ΔG0 is large and negative; when Keq