Theoretical Fundamentals of Bioenergetics

1. Theoretical Fundamentals of Bioenergetics

2. THERMODYNAMICS– is the branch of science which deals with energy changes in physical and chemical processes.

3. Terms used in Thermodynamics System - is a specified part of the universe under observation. Surroundings- the remaining portion of the universe which is not a part of the system. The system and surroundings together are called “universe”.

4. Types of Systems Isolated system-is a system which can exchange neither mass nor energy with the surroundings. For example, chemical reactions carried out in closed containers insulated from the surroundings (Δm=0; ΔE=0). Closed system-is a system which can exchange energy but not mass with its surroundings. For example, liquid in a closed container (Δm=0; ΔE≠0). Open system-is a system which can exchange matter as well as energy with the surroundings (Δm ≠0; ΔE≠0). Biological systems, or living  organisms, are considered open thermodynamic systems.

5. Types of Systems Homogeneous system – has single or uniform phase in the form of solid, liquid or gas. Heterogeneous system - has more than one phase i.e. the combination of solid, liquid and gaseous state.

6. Parameters of a System Parameters- physical and chemical properties which characterize the state of a system. P, T, V, m, C The state of a system–is the condition of existence of the system when its parameters have definite values. If any of the parameters of the system changes, the state of the system is also changes.

7. State Functions Energetic state of a system is characterized by state functions. A state function is a property of the system whose value depends only upon the state of the system and is independent of the path or manner by which the state is reached. The change in the value of state functions depends only upon the initial and final states of the system and not on the path by which the change from initial to final state is brought about. Internal energy (U), Enthalpy (H), Entropy (S), Gibbs free energy (G)

8. Types of processes Isothermal process: The process during which the temperature of the system remains constant. (T=const.) Isochoric process: In this process the volume of system remains constant. (V=const.) Isobaric process: The process during which the pressure of the system remains constant. (P=const.)

9. First Law of Thermodynamics Energy can be neither created nor destroyed, it can be converted from one form into another. The energy can neither be created nor destroyed but internal energy of the system can be changed in two ways: 1. by allowing heat to flow into the system or out of the system. 2. by work done on the system or by work done by the system.

10. Mathematical statement of the First Law of Thermodynamics Q = ΔU + A V=constQ=ΔU. Heat absorbed by the system in isochoric process is equal to the change in internal energy. Therefore, the total energy of the universe is a constant. Energy can, however, be converted from one form to another or transferred from a system to surroundings or vise versa.

11. First Law of Thermodynamics

12. Chemical reactions are generally carried out at a constantpressure, i.e. at atmospheric pressure. at P=const,Q=ΔU+A ΔH Q = ΔH Enthalpy change (ΔH) is a measure of heat change (evolved orabsorbed) taking place during a process at constant pressure.

13. Heat Effect of the Reaction In a chemical reaction, the change in the enthalpy equals the heat effect of the reaction conducted at T = const. and P = const. taken with the opposite sign (Q = -ΔH). exothermic reactionΔH <0 heat released endothermic reactionΔH >0 heat absorbed Enthalpy of the reaction – is the enthalpy change of a balanced reaction. CH4(g) + 2O2(g)→CO2(g) + 2H2O(g)ΔH = -890.3kJ

14. Enthalpy Changes Endothermicreaction: ΔHrxn = positive Heat must be added to the system The decomposition of water is endothermic! Q = - ΔH P H R time 2 H2O (g)  2 H2 (g) + O2 (g) DH = + 483.6 kJ

15. Enthalpy Changes Exothermic reaction: ΔHrxn = negative Heat is lost by the system Combustion of methane is exothermic. Q = - ΔH R H P time CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (l) DH = -890. kJ

16. Standard-state conditions The standard enthalpy of a reaction (ΔH°) - is the value of enthalpy change when the reaction is carried out at standard conditions (1 atm; 298K). The enthalpy of formation (ΔHf) is defined as the enthalpy change when one mole of a compound is formed from its elements. C(s) + O2(g)→CO2(g) ΔH = ΔHf = -393.5kJ If all the species of the chemical reaction are in their standard state, (i.e. at 1 atm pressure and 298 K) the enthalpy of formation is called standard enthalpy of formation (ΔHf).

17. Standard Enthalpy Standard enthalpy of combustion (ΔH°c) of a substance is the amount of heat evolved when 1 mole of a substance is completely burnt or oxidized under standard conditions. The enthalpy of formation of every element in its standard state is assumed to be zero, because there is no formation reaction needed when the element is already in its standard state.

18. Hess’ Law Q3 Q4 Q2 Q1 A B Q5 Q6 The heat effect (enthalpy) of a reaction depends only on the initial and the final state of substances and does not depend on the intermediate stages of the process. Q1 = Q2+Q3+Q4 = Q5+Q6

19. Corollaries of Hess`s Law 1. ΔH°reaction=Σ ΔH°formation of products - Σ ΔH°formation of reactants 2. ΔH°reaction=Σ ΔH°combustion of reactants - Σ ΔH°combustion of products Path 1.C(s) + O2(g)→CO2(g)Δ H1= -393.5kJ Path 2.C(s) + 1/2O2(g)→CO(g)ΔH2= -110.5kJ CO(s) + 1/2O2(g)→CO2(g)ΔH3= -283.5kJ __________________________________________________ C(s) + O2(g)→CO2(g)ΔH2=-283+(-110.5)= -393.5kJ ΔH1 = ΔH2 + ΔH3 ΔH2 = ΔH1 – ΔH3 We can NOT predict whether a reaction will occur based only on ∆H!

20. First Law for Living Organisms Organism does not produce new energy, it works on account of energy stored in nutrients!

21. Energy expenditure in humans

22. Direction of a process The first law of thermodynamics does not determine the feasibility a process. This question is decided bySecond law of Thermodynamics. Spontaneous processes are those that can proceed without any outside intervention The spontaneous processes occur because of the two tendencies: 1.Tendency of a system to acquire a state of minimum energy 2. Tendency of a system to acquire a state of maximum randomness Entropy (S), J/mol∙K - the extent of disorder or randomness in a system.

23. Spontaneous Processes • Processes that are spontaneous in one direction are nonspontaneous in the reverse direction. • Processes that are spontaneous at one temperature may be not spontaneous at other temperature. Above 0oC it is spontaneous for ice to melt. • Below 0oC it the reverse process is spontaneous. • At 0oC the reactions are in equilibrium.

24. Second Law of Thermodynamics The entropy always increases in the course of every spontaneous (natural) change.

25. Entropy Entropy and Physical States Entropy increases with the freedom of motion of molecules. S gas > S liquid > S solid The entropy tends to increase with increases in • Temperature. • Volume. • Decrease of pressure • The number of independently moving molecules.

26. Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is zero. Entropies of pure substances can be determined by experiment. Standard entropy (S°) – is the entropy of one mole of the substance at 298K and 1 atm pressure. ΔS°(reaction)=ΣS°(products) - ΣS°(reactants)

27. Gibbs Helmholtz Equation ΔG =Δ H - TΔS Where, ΔG – is Gibbs free energy change . Gibbs free energy is the maximum amount of energy available to a system during a process that can be converted into useful work. The free energy change, ΔG provides the overall criterion for the feasibility of a chemical process.

28. Conditions for Spontaneous Process ΔG < 0 – process is spontaneous ΔG = 0 – process in equilibrium ΔG > 0 – process does not occur by itself ΔG =Δ H - TΔS ΔH<0, ΔS>0–process is spontaneous at any temperature ΔH>0, ΔS<0– process is non-spontaneous ΔH<0, ΔS<0– process is spontaneous at low temperature ΔH>0, ΔS>0– process is spontaneous at high temperature ΔG°=Σ Δ Gf°(products) - Σ Δ Gf°(reactants)

29. Exergonic and Endergonic Reactions Exergonic reactions – biochemical reactions in which Gibbs free energy decreases. (ΔG<0) Endergonic reactions - biochemical reactions in which Gibbs free energy increases. (ΔG>0)

30. Homeostasis Homeostasis - The ability or tendency of an organism or a cell to maintain internal equilibrium by adjusting its physiological processes. Cells make endergonic reaction happen by supplying them with free energy released by exergonic reactions. Energy coupling: The transfer of energy in order to drive the second reaction. In living cells the exergonic reaction is the hydrolysis of the ATP to the ADP. ATP + H2O → ADP + HPO42-Δ G°= -7.3 kcal/mole (-30,4 kJ/mole) ATP is a high-energy (macroergic) compound.

31. Concept of coupling reactions

32. Living Systems • Living organisms are open thermodynamic systems. • Biological systems continuously exchange mass as well as energy with surroundings • Processes in biological systems are irreversible • The state of living system can be characterized as a stable non-equilibrium stationary state – homeostasis (constant chemical composition of internal medium, osmotic pressure, pH, temperature, etc).