- BIOENERGETICS - DR. A. TARAB DEPT. OF BIOCHEMISTRY HKMU

Transcription

1 - BIOENERGETICS - DR. A. TARAB DEPT. OF BIOCHEMISTRY HKMU

2 Bioenergetics the field of biochemistry concerned with the transfer and use of energy by biological system BIOLOGICAL IMPORTANCE: Suitable fuel is required to provide the energy that enables the animal to carry out its normal processes How the organism obtains this energy from its food is basic to the understanding of normal nutrition and metabolism

3 Death from starvation occurs when available energy reserves are depleted, and certain forms of malnutrition are associated with energy imbalance (kwashiorkor, marasmus) Storage of surplus energy results in obesity Of key importance in bioenergetics is the application of basic principles of thermodynamics (branch of physical science that deals with energy transformations) to the study of energy transformations within cells

4 Energy is a highly abstract concept, which is best understood and measured by its effects Energy is defined as the capacity to do work, and it is denoted by the symbol E It can manifest itself in different forms such as mechanical, thermal, chemical, electrical or radiant energy

5 Chemical energy the energy inherent in the structure, bonds and configuration of molecules It may be released or absorbed as a result of structural rearrangements and chemical transformations If the total energy of the products of a reaction is lower than that of the reactants, the difference in energy ΔE, is released to the environment Conversely, when it is desired to generate products with higher energy than those of the reactants, additional energy ΔE, has to be imparted into the system

6 Thermodynamic systems and their surroundings The collection of matter (e.g. a chemical reaction, a cell or an entire organism) in which we wish to study the energy changes during some chemical or physical process is called a system All the other matter in the universe apart from this system, is called the surroundings

7 If a system and its surroundings can exchange both matter and energy, the system is defined as open If energy, but no matter can be exchanged, the system is defined as closed If neither energy nor matter can be exchanged between system and surroundings, the system is considered isolated Biochemical experiments can be designed and executed under any of these conditions

8 State functions (thermodynamic variables): - When the state of a system changes, the new system state is determined by the difference between the initial and final values of the state functions, independent of the specific pathway of the process by which the system reached its new state A variable that is dependant on the pathway, such as work done, is not a state function

9 Energy, temperature, pressure and component concentrations are examples of state functions that can be analyzed to determine changes in thermodynamic state

10 The first law of thermodynamics The first law, enunciated by Robert Mayer in 1841, is the principle of the conservation of energy energy can be neither created nor destroyed This, in any chemical or physical process, the total energy of the system plus surroundings, that is, the total energy of the universe, remains constant In terms of a single chemical reaction: Energy + aa + bb cc + dd

11 where a, b, c and d are the number of molecules of A, B, C and D respectively The energy absorbed in the forward reaction is exactly equal to the energy released in the reverse reaction The energy change in reaction is given by ΔE = Q W where ΔE change in internal energy Q heat absorbed by the system W work done by the system

12 The quantity of ΔE is the same for conversion of A and B to C and D whether the process requires many steps and involves many intermediates, as it well might in a metabolic pathway, or whether, it occurs in a single step The work done by a biological system includes, for example, the mechanical work of muscle contraction, the electrical work of nerve transmission and work done in the transport of substances across cell membranes

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14 Almost all biological processes occur at constant pressure One form of work (W) done by a system at constant pressure equals the product of the pressure (P) and the change in volume (ΔV) that occurs during the process W = PΔV Substituting and rearranging, we obtain Qp = ΔE + PΔV

15 where Qp is the heat consumed in the process at constant pressure which is also a state function The heat released or absorbed in a process occurring at constant pressure is the enthalpy change (ΔH) for the reaction Thus, equation can be rewritten ΔH = ΔE + PΔV Enthalpy change (ΔH) is also a state function

16 The second law of thermodynamics It states that - the total disorder of the universe increases in every process ΔS process = ΔS system + ΔS surroundings where S is the symbol for entropy (extent of disorder or randomness of the system A system is at equilibrium with its surroundings when entropy has increased to a maximum, and no energy is available to drive processes

17 Thus, the tendency toward increasing entropy acts as a force that drives systems towards equilibrium with their surroundings When an organism is at equilibrium with its surroundings, it is dead

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19 The concept of free energy unifies the first and second laws of thermodynamics The ability of a system to do work decreases as equilibrium is approached, and at equilibrium there is no energy available to do work Gibbs free energy (G) is a thermodynamic state function that defines the equilibrium condition in terms of the enthalpy and entropy of the system at a constant temperature and pressure ΔG = ΔH - TΔS

20 where ΔG is the Gibbs free energy change T is the Kelvin temperature Since the enthalpy change is a property of the first law of thermodynamics and the entropy change is a property of the second law, the Gibbs free energy change provides a unifying principle in thermodynamics The equation was formulated in 1878 by Josiah Willard Gibbs, who combined the two laws of thermodynamics

21 The Gibbs free energy change can be positive, negative or zero 1) If ΔG is negative in sign, the reaction proceeds spontaneously with loss of free energy; it is exergonic If, in addition ΔG is of great magnitude, the reaction goes virtually to completion and is essentially irreversible

22 2) If ΔG is positive, the reaction proceeds only if free energy can be gained, i.e., it is endergonic If, in addition, the magnitude of ΔG is great, the system is stable with little or no tendency for a reaction to occur 3) If ΔG is zero, the system is at equilibrium with no net change takes place

23 Endothermic and exothermic reactions In thermodynamics, the term endothermic describes a process or reaction in which the system absorbs energy from its surroundings in the form of heat The term was coined by Marcellin Berthelot from the Greek roots endo-, derived from the word endon (ἔνδον) meaning within and the root therm (θερμ-) meaning hot

24 The intended sense is that of a reaction that depends on taking in heat if it is to proceed The opposite of an endothermic process is an exothermic process, one that releases, gives out energy in the form of heat Thus in each term (endothermic and exothermic) the prefix refers to where heat goes as the reaction occurs

25 Examples of endothermic reactions: - photosynthesis, melting ice, thermal decomposition reactions, etc. Examples of exothermic reactions: - combustion reactions of fuels, burning of a substance, neutralization reactions, etc.

26 An exothermic thermite reaction using iron(iii) oxide. The sparks flying outwards are globules of molten iron trailing smoke in their wake *Thermite: a mixture of aluminium powder and a metal oxide, such as iron oxide

27 The concept is frequently applied in physical sciences to, for example, chemical reactions, where thermal energy (heat) is converted to chemical bond energy

28 The term exergonic and endergonic rather than the normal chemical terms exothermic and endothermic are used to indicate that a process is accompanied by loss or gain, respectively of free energy in any form, not necessarily as heat

29 Biological thermodynamic standard state When the reactants are present in concentrations of 1.0 mol/l, ΔG o is the standard free energy change For biochemical reactions, a standard state is defined as having a ph of 7.0 The standard free energy change at this standard state is denoted by ΔG o

30 The standard free energy change is related to the equilibrium constant for a reaction: ΔG o = RT log k eq where R - is the gas constant (1.987 calories per mole per degree) T - is the absolute temperature (degrees Kelvin, or degrees Celsius plus 273)

31 Energy changes in metabolic reactions are often coupled The secret of energy flow in biological systems is that energy released by one process is coupled to a second process that would not otherwise occur spontaneously under cellular conditions Thus the energy produced by spontaneous processes is not wasted but instead provides the thermodynamic driving force for nonspontaneous processes

32 Coupled biochemical reactions are catalyzed by enzymes An enzyme does not change the equilibrium constant for a reaction and does not provide the driving force for the reaction; rather, it provides a pathway that enables the reaction to occur The simplest case of energy coupling occurs when two metabolic reactions share a common intermediate

33 k 1 A + B C ΔG o 1 k 2 C + D E ΔG o 2 The total standard free energy change for the process (ΔG o 3) is the sum of the standard free energy changes for the individual steps k 1 xk 2 =k 3 A + B + D E ΔG o 3 = ΔG o 1 + ΔG o 2

34 The conversion of malate to aspartate illustrates the general principle of coupled reactions that share a common intermediate First, fumarase converts malate to fumarate, then aspartase converts fumarate to aspartate

35 malate fumarate + H 2 O ΔG o 1 = kcal mol -1 k 1 = 0.21 NH 4+ + fumarate aspartate ΔG o 2 = kcal mol -1 k 2 = 5.28 x 10 2 The sum of the two reactions is NH 4+ + malate aspartate + H 2 O ΔG o 3 = kcal mol -1 K 3 = 1.11 x 10 2

36 Hexokinase reaction

37 D-Glucose + Pi G6P + H 2 O ΔG o = kcal mol -1 ATP + H 2 O ADP + Pi ΔG o = kcal mol -1 D-Glucose + ATP G6P + ADP ΔG o = kcal mol -1

38 ATP mediates biological energy transfer The flow of energy in metabolism revolves around coupled reactions in which ATP is the intermediate Most of the coupled reactions in which ATP participates involve phosphoryl-group transfer from ATP to another substance of from an energy-rich metabolite to ADP, forming ATP

39 The standard free energy of the hydrolysis of energy rich metabolites ranges from 7 to 14 kcal mol -1 The standard free energy of hydrolysis of ATP is about 8 kcal mol -1 ATP is not the only nucleoside triphosphate to serve in this capacity, and we shall encounter transformations in which uridine triphosphate (UTP), cytidine triphosphate (CTP) or guanosine triphosphate (GTP) participate

40 Adenosine triphosphate (ATP) The energy released by the degradation of nutrients is converted to a specific intermediate form of chemical energy, which the cells can channel for the metabolic and mechanical activities of life This central energy intermediate is ATP It is unique in that the cellular machinery is geared to its formation when energy is released and to its utilization when energy is required

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42 The energy-releasing reaction of ATP is its hydrolysis, which yields adenosine diphosphate (ADP) and inorganic phosphate (Pi) If pyrophosphate group (PPi) is removed adenosine monophosphate (AMP) results The bonds between the phosphates are anhydride bonds, whereas that between phosphate and carbon-5 of ribose is an ester bond

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44 Hydrolysis of any anhydride is a highly exergonic process Consequently, hydrolytic cleavage of either the terminal or the second phosphate of ATP has a high, negative ΔG In contrast, the hydrolysis of the remaining phosphate from AMP does not result in the release of exceedingly high amounts of energy (ΔG o = kilocalories per mole)

45 In ATP, there are four neighboring negative charges that repel each other; this creates an electrical tension or highenergy state in the molecule ATP has a high affinity for magnesium ions, and within the cell, it exists primarily as a magnesium complex

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47 This interaction with magnesium and other intracellular ions also has a critical effect on the free energy of hydrolysis

48 Hydrolysis of energy-rich metabolites can be coupled with synthesis of ATP The hydrolysis of some other energy-rich metabolites releases more free energy than the hydrolysis of ATP These metabolites can transfer a phosphoryl group to ADP, forming ATP These coupled reactions provide ATP for a variety of cellular processes ranging from biosynthesis to cell movement

49 Carbamyl phosphate, an intermediate in the urea cycle and in pyrimidine biosynthesis, is a mixed anhydride of carbamic acid and phosphoric acid The standard free energy of hydrolysis for carbamyl phosphate is 11.7 kcal mol-1

50 1,3-bisphosphoglycerate (1,3 BPG) is an intermediate in glycolysis The standard free energy of hydrolysis of 1,3BPG to 3-phosphoglycerate (3PG) is 11.8 kcal mol -1 1,3BPG + ADP 3PG + ATP phosphoglycerate kinase

51 The concentration of creatine phosphate in muscle is about 20 mm, 10 times that of ATP Creatine phosphate thus provides an immediate and local reservoir of high energy phosphate for muscle contraction creatine phosphate + ADP creatine + ATP creatine phosphokinase ΔG o hydrolysis = kcal mol -1

52 Creatine phosphokinase reaction

53 Phosphoenolpyruvate (PEP) is an intermediate in glycolysis whose standard free energy of hydrolysis is 14.8 kcal mol -1 PEP + ADP + H + pyruvate + ATP pyruvate kinase

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55 Units of energy A calorie (cal) is equivalent to the amount of heat required to raise the temperature of 1 gram of water from 14.5 C to 15.5 C. A kilocalorie (kcal) is equal to 1000 cal. A joule (J) is the amount of energy needed to apply a 1-newton force over a distance of 1 meter. A kilojoule (kj) is equal to 1000 J. 1 kcal = kj