BIOENERGETICS - Copy.ppt

BIOENERGETICS  Living organisms must work to stay alive, to grow and to reproduce All living organisms have the ability to produce energy and to channel it into biological work Living organisms carry out energy transductions, conversions of one form of energy to another form  Mod...

BIOENERGETICS  Living organisms must work to stay alive, to grow and to reproduce All living organisms have the ability to produce energy and to channel it into biological work Living organisms carry out energy transductions, conversions of one form of energy to another form  Modern organisms use the chemical energy in fuels (carbonhydrates, lipids) to bring about the synthesis of complex macromolecules from simple precursors They also convert the chemical energy into concentration gradients and electrical gradients, into motion and heat, and, in a few organisms into light (fireflies, some deep-sea fishes)  Biological energy transductions obey the same physical laws that govern all other natural processes  Bioenergetics is the quantitative study of the energy transductions that occur in living cells and of the nature and function of the chemical process underlying these transductions  BIOENERGETICS & THERMODYNAMICS Biological energy transductions obey the laws of thermodynamics  Laws of thermodynamics 1.For any physical or chemical change, the total amount of energy in the universe remains constant. Energy may changed from or it may be transported from one region to another, but it can not be created or destroyed 2.The universe always tends toward increasing disorder: In all natural processes the entropy of the universe increases  The reacting system 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 closed systems and no material or energy is exchanged with the surroundings However living organisms are open systems.They exchange both material and energy with their surroundings  Living systems are never at equilibrium with their surroundings Gibbs free energy (G): G 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 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 The unit of ΔG is joules/mole or calories/mole  Enthalpy (H): H is the heat content of the reacting system. 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 ΔH has a negative value Reacting systems that take up heat from their surroundings are endothermic and have positive values of ΔH  The unit of ΔH is joules/mole or calories/mole Entropy (S): S is a quantitative expression for the randomness or a disorder in a system The unit of ΔS is joules/mole. Kelvin  Under the constant temperature and pressure changes in free energy, enthalpy and entropy in biological systems are related to each other by the equation ΔG= ΔH - TΔS ΔG= Change in Gibbs free energy of the reacting system (Gproducts– Greactives) ΔH= Change in enthalpy of the reacting system(Hproducts – Hreactives) T= Absolute temperature ΔS= Change in entropy of the reacting system(Sproducts – Sreactives) Living organisms preserve their internal order by taking free energy from their surroundings in the form of nutrients or sunlight and returning to their surroundings an equal amount of energy as heat and entropy  Example:The oxidation of glucose Aerobic organisms extract free energy from glucose obtained from their surrouindings by oxidizing the glucose with oxygen (also obtained from surroundings). The end products of this oxidation reaction are CO2 and H2O and they are returned to the surroundings. At the end of this process, the surroundings undergo an increase in entropy, whereas the organism itself remains in a steady state and no change occurs in its internal order  Whenever a chemical reaction results in an increase in the number of molecules the entropy of the surroundings increases Whenever a solid substance is converted into liquid or gaseous forms the entropy of the surroundings increases. Because this transformations allow the molecule more freedom for movement Cells require sources of free energy Living organisms acquire free energy from nutrient molecules. Cells transform this free energy into ATP and other energy-rich compounds They are capable of providing energy for biological work at constant temparature.  The composition of a reacting system tends to continue changing until equilibrium is reached. At the equilibrium the rates of the forward and revers reactions are equal and no further change occurs in the system.  The Keq is defined by the molar concentrations of products and reactants at equilibrium aA+bB cC+ dD [C]c [D]d Keq = [A]a [B]b  When a reacting system is not at equilibrium, the tendency to move toward the equilibrium represents a driving force. The magnitude of this driving force is expressed as free energy change (ΔG).  ΔG'0 is the difference between the free energy content of the products and the free energy contents of the reactants under standard conditions ΔG'0 = ΔG'0 products– ΔG'0 reactives) When ΔG'0 is negative, the products contain less free energy than the reactants and the reaction will proceed spontaneously under standard conditions When ΔG'0 is positive, the products contain more free energy than the reactants and the reaction will tend to go in the revers direction under standard conditions  Each chemical reaction has a characteristic standard free energy change which may be positive, negative or zero depending on the equilibrium constant of the reaction ΔG'0 tell us in which direction and how far a given reaction must go to reach equilibrium when the initial concentration of each component is 1M, the pH is 7, the temparature is 250C. Thus ΔG'0 is a constant; a characteristic for a given reaction  Standard free energy changes are additive. In the case of two sequential chemical reactions, A B ΔG'01 B C ΔG'02  Since the two reactions are sequential, we can write the overall reaction as A C ΔG'0total The ΔG'0 values of sequential reactions are additive. ΔG'0total = ΔG'01 + ΔG'02  A B ΔG'01 B C ΔG'02 Sum: A C ΔG'01 + ΔG'02 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 through a common intermediate  The main rule in biochemical reactions in living organisms: All endergonic reactions are coupled to an exergonic reaction. There is an energy cycle in cells that links anabolic and catabolic reactions.  An example: The first step of glycolysis Glucose + Pi Glucose 6-phosphate+H2O ΔG'0 =13.8 kj/mol ΔG'0 >0 reaction is not spontaneous  Another very exergonic cellular reaction: Hydrolysis of ATP ATP + H2O ADP + Pi ΔG'0 = -30.5 kj/mol These two reactions share the common intermediates H2O and Pi and may be expressed as sequential reactions:  Glucose + Pi Glucose 6-phosphate+H2O ATP + H2O ADP + Pi Sum: Glucose+ATP Glucose 6-phosphate+ADP The overall standard free energy changes: ΔG'0 =13,8 kj/mol + (-30,5 kj/mol)= -16.7 kj/mol Overall reaction is exergonic Energy stored in ATP is used to drive to synthesis of glycose 6-phosphate, eventhough its formation from glucose and Pi is endergonic. This strategy works only if compounds such as ATP are continuously available.  Phosphoryl group transfer and ATP Living cells obtain free energy in a chemical form by the catabolism of nutrient molecules They use that energy to make ATP from ADP and Pi. ATP donates some of its chemical energy to: 1. Endergonic processes such as the synthesis of metabolic intermediates and macromolecules from smaller precursors 2.The transport of substances across membranes against concentration gradients 3.Mechanical motion (muscle contraction) This donation of energy from ATP can occur in the two form A) ATP ADP+ Pi or B) ATP AMP+ 2 Pi  The free energy change for ATP hydrolysis is large and negative The hydrolytic cleavage of the terminal phosphoanhydride bond in ATP separates one of the three negatively charged phosphates and thus relieves some of the electrostatic repulsion in ATP Released Pi is stabilized by the formation of several resonance forms not possible in ATP  Although the hydrolysis of ATP is highly exergonic (ΔG'0 = -30.5 kj/mol), the ATP is stable at pH 7, because the activation energy for ATP hydrolysis is relatively high. Rapid hydrolysis of ATP occurs only when catalyzed by an enzyme  Mg2+ in the cytosol binds to ATP and ADP and for most enzymatic reactions that involve ATP as phosphorly group donor, the true substrate is MgATP- 2. The relevant ΔG'0 is therefore that for MgATP-2 hydrolysis.  Compounds have large free energy change Phosphorylated compounds Thioesters (Acetyl-CoA)  Phosphorylated compounds Phosphoenolpyruvate 1,3-bisphosphoglycerate Phosphocreatine ADP ATP AMP PPi Glucose 1-phosphate Fructose 6-phosphate Glucose 6-phosphate  Summary for hydrolysis reactions For hydrolysis reactions with large, negative standard free energy changes, the products are more stable than the reactants for one or more of the following reasons: 1.The bond strain in reactants due to electrostatic repulsion is relieved by charge separation, as for ATP 2.The products are stabilized by ionization, as for ATP, acyl phosphates, thioesters 3. The products are stabilized by isomerization (tautomerization) as for phosphoenolpyruvate 4. The products are stabilized by resonance as for creatine released from phosphocreatine, carboxylate ion released from acyl phosphates and thioesters and phosphate released from anhydride or ester linkages  The phosphate compounds found in living organisms can be arbitrarily divided into two groups based on their standard free energy changes of hydrolysis. 1. High-energy’ compounds have a ΔG'0 of hydrolysis more negative than -25 kj/mol; ‘low-energy’ compounds have a less negative ΔG'0. 2. Based on this criterion, ATP, with a ΔG'0 of hydrolysis of -30 kj/mol is a high-energy compound; glucose 6-phosphate is a low-energy compound (ΔG'0 = -13,8 kj/mol)  One more chemical feature of ATP is crucial to its role in metabolism: although in aqueous solution ATP is thermo- dynamically unstable and is therefore a good phosphoryl group donor, it is kinetically stable. Because of high activation energies required for uncatalyzed reaction ATP does not spontaneously donate phosphoryl groups to water or to the other potential acceptors in the cell.  ATP hydrolysis occurs only when specific enzymes which lower the energy of activation are present The cell is therefore able to regulate the disposition of the energy carried by ATP by regulating the various enzymes that act on ATP Each of the three phosphates of ATP is susceptible to nucleophilic attack and each position of attack yields a different type of product  Attack at the Ɣ phosphate displaces ADP and transfers Pi, Attack at the ϐ phosphate displaces AMP and transfers PPi, Attack at the α phosphate displaces PPi and transfers adenylate Notice that hydrolysis of phospho anhydride bond between α and ϐ phosphates releases much more energy than hydrolysis of phospho anhydride bond between ϐ and Ɣ phosphates. Because PPi formed as a by-product of the adenylation is hydrolyzed the two Pi releasing and thereby providing a further energy push for the adenylation reaction  SUMMARY ATP is the chemical link between catabolism and anabolism. The exergonic conversion of ATP coupled to many endergonic processes in living organisms Cells also contain some high-energy compounds which have a high phosphorylation potential, like ATP. They are good donors of phosphoryl groups.