1. THEME: Theoretic bases of bioenergetics. Chemical kinetics and biological processes. Electrochemistry. associate. prof. Yevheniy. B. Dmukhalska LECTURE № 3

2. Plan 1.The basic concepts of thermodynamics 2. First law of thermodynamics. Heat (Q) and Work ( W) 3. Secohd law of thermodynamics. Entropy (S) 4. Electrochemistry.

3. 3 The branch of science which deals with energy changes in physical and chemical processes is called thermodynamics Some common terms which are frequently used in the discussion of thermodynamics are:

4. 4 Common terms of thermodynamicsSystem Parameter Condition (state) Process characterized characterizes

6. 6 System System is a specified part of the universe which is under observation surroundings The remaining portion of the universe which is not a part of the system is called the surroundings The system is separated by real or imaginary boundaries.

7. Classification of the thermodynamics systems according to a structure homogeneous heterogeneous KNO 3 PbI 2 ↓ KNO 3

8. 8 Types of Systems OpenClose ISOLATED A system can neither exchange matter nor energy with the surroundings CLOSE A system which can exchange energy but no mass with its surroundings OPEN A system can exchange both matter and energy with the surroundings.

9. 9 Parameters Extensive (m, V, U, H, G, S, c) The properties of the system whose value depends upon the amount of substance present in the system Intensive (p, T, C, viscosity, surface tension, vapour pressure) The properties of the system whose value does not depend upon the amount of substance present in the system

10. 10 Process Process is the change of all or individual parameters of the system during the length of time (the period of time) Classification of a process according to the constant parameter of a system are: Isothermic process – temperature is constant, T=const Isochoric process – volume is constant V = const. Isobaric process – pressure of the system is constant, p = const Adiabatic process – the system is completely isolated from the surroundings. For an adiabatic (Q=0) system of constant mass, ▲U=W

11. 11 Classification of a process according to the releasing energy Exothermic process is a process that releases energy as heat into its surroundings. We say that in an exothermic process energy is transferred ‘as heat’ to the surroundings. For example: a reaction of neutralization (acid + basic). Endothermic process is a process in which energy is acquired from its surroundings as heat. Energy is transferred ‘as heat’ from the surroundings into the system. For example: the vaporization of water

12. 12 Classification of a process according to the direction of reaction. Reversible process is a process in which the direction may be reversed at any stage by merely a small change in a variable like temperature, pressure, etc. Irreversible process is a process which is not reversible. All natural process are irreversible

13. 13 State of a system means the condition of the system, which is described in terms of certain observable (measurable) properties such as temperature (T), pressure (p), volume (V) State function (thermodynamic function) Internal energy U [J/mol] Enthalpy H [kJ/mol] or [kJ] Entropy S [J/mol K] or [J/K] Gibbs energy G [J/mol] or [J] ΔU = U(products) – U(reactants)

14. 14 State function depends only upon the initial and final state of the system and not on the path by which the change from initial to final state is brought about.

15. 15 Internal energy U It is the sum of different types of energies associated with atoms and molecules such as electronic energy, nuclear energy, chemical bond energy and all type of the internal energy except potential and kinetic energies.

16. 16 Heat (Q) is a form of energy which the system can exchange with the surroundings. If they are at different temperatures, the heat flows from higher temperature to lower temperature. Heat is expressed as Q.

17. 17 Work (W) is said to be performed if the point of application of force is displaced in the direction of the force. It is equal to the distance through which the force acts.

18. 18 Enthalpy H Chemical reactions are generally carried out at constant pressure. ΔU gives the change in internal energy at constant volume. To express the energy changes at constant pressure, a new term called enthalpy was used. Enthalpy cannot be directly measured, but changes in it can be.

19. 19 Enthalpy H A thermodynamic function of a system, equivalent to the sum of the internal energy of the system plus the product of its volume multiplied by the pressure exerted on it by its surroundings. ▲H = ▲U + p▲V

20. 20 The meaning of the state functions in the thermodynamic processes process Exothermic process Qv > 0, ▲U Qv > 0, ▲U < 0 Qp > 0, ▲H Qp > 0, ▲H < 0 Endothermic process Qv 0, ▲U> Qv 0 Qp 0, ▲H> Qp 0 Heat absorbed by the system = H positive (Q negative Heat evolved by the system = H negative (Q positive) The signs of W or Q are related to the internal energy change.

21. 21 The first law of thermodynamics Matter/energy may be altered (converted), but not created (from nothingness) nor destroyed (reduced to nothingness). The First Law teaches that matter/energy cannot spring forth from nothing without cause, nor can it simply vanish. Energy can neither be created nor destroyed although it may be converted from one form to another. The given heat for the system spends on the change of the internal energy and producing the work: Q = ▲U + W

22. 22 Bomb calorimeter for the determination of change in internal energy The process is carried out at constant volume, i.e., ΔV=0, then the product PΔV is also zero. Thus, ΔU=Qv The subscript v in Qv denotes that volume is kept constant. Thus, the change in internal energy is equal to heat absorbed or evolved at constant temperature and constant volume

23. 23 Thermochemistry The study of the energy transferred as heat during the course of chemical reactions. Thermochemical reactions: H 2 (g) + Cl 2 (g) = 2HCl; ▲ H = -184,6 kJ 1/ 2 H 2 (g) + 1/ 2 Cl 2 (g) = HCl; ▲ H = -92,3 kJ/mol ▲ H is calculated for 1 mole of product ▲H = ▲U + p▲V ▲H = ▲U + ▲nRT Energy change at constant P = Energy change at constant V + Change in the number of geseous moles * RT

24. 24 If the volume or pressure are constant the total amount of evolved or absorbed heat depends only on the nature of the initial reactants and the final products and doesn’t depend on the passing way of reaction. The Hess’s law Н1Н1 Н2Н2 Н3Н3 Н4Н4 Initial reactants The products of reaction  Н 1 =  Н 2 +  Н 3 +  Н 4

25. 25 Conclusions from the Hess law 4.  Н 3 =  Н 1 -  Н 2 1 2 3 5.  Н 1 =  Н 3 -  Н 2 1 2 3 1.  Н c 298 (the standard enthalpy of combustion) =-  Н f 298 (the standard enthalpy of formation) 2.  Н( formation) = Σn  Н f 298 (products) - Σn  Н f 298 (reactants) 3.  Н( combustion) = Σn  Н с 298 (reactants) - Σn  Н с 298 (products) 4.For elementary substances  Н 0 298 = 0

26. 26 Correlation  U і  Н: If  υ  0, so  Н  U: СаО + СО 2 → СаСО 3 If  υ  0, so  Н  U: Na + H 2 O → NaOH + H 2 If  υ=0, so  Н =  U: H 2 + Cl 2 → 2HCl

27. Second law of thermodynamics Second Law of Thermodynamics (refrigerator): It is not possible for heat to flow from a colder body to a warmer body without any work having been done to accomplish this flow. Second Law of Thermodynamics (refrigerator): It is not possible for heat to flow from a colder body to a warmer body without any work having been done to accomplish this flow.

28. The amount of molecular randomness in a system is called the system’s entropy (S). Entropy is a measure of randomness or disorder of the system

29. 29 Free energy and free energy change The maximum amount of energy available to a system during a process that can be converted into useful work It’s denoted by symbol G and is given by ▲ G = ▲ H - T ▲ S where ▲ G is the change of Gibbs energy (free energy) This equation is called Gibbs equation and is very useful in predicting the spontaneity of a process. N.B. Gibbs equation exists at constant temperature and pressure

30. 30 1) Spontaneous (irreversible) process : ▲ G 0, ▲ H 0, ▲ H < 0 2) Unspontaneous (reversible) process : ▲ G > 0, ▲ S 0 3) Equilibrium state ▲ G = 0

31. 31 THIRD LAW OF THERMODYNAMICS: The third law of thermodynamics, formulated by Walter Nernst and also known as the Nernst heat theorem, states that if one could reach absolute zero, all bodies would have the same entropy.

32. A chemical kinetics of biological processes

33. Definition Chemical kinetics is that branch of chemistry, which deals with the study of the the rates of chemical reactions, the factors affecting the rates of the reactions and the mechanism by which the reactions proceed.

34. 34 Classification of chemical reactions according to the quantity of stages (phases). Simple reactions go in a one elementary chemical act Complex(compound) reactions go in several stages

35. Chain reactions Primary process – chain initiating step (stage): h  Cl 2 === 2С1. chlorine molecule absorbs one quantum of light (h  ) and dissociates to give Cl atoms. Secondary process – chain propagating step (stage): 1. Cl. + Н 2 = HCl + H. 2. H. + Cl 2 = HCl + Cl. Third process – chain terminating step (stage): Сl. + Cl. = Сl 2

36. Parallel reactin For example: Phenol with nitric acid, so have been formed ortho-, pair- and meta-nitrophenol. Series the reactions are reaction which products firs step ( stage ) are reactants for second step ( stage ): A  B  C  D  …. C 18 H 32 O 16 + HOH = C 12 H 22 O 11 + C 6 H 12 O 6 Raffinose disaccharide monosaccharide C 12 H 22 O 11 + HOH = C 6 H 12 O 6 + C 6 H 12 O 6 Monosaccharides

37. Reversible the reactions reactions which are flowing past in two parties: the forward reaction - conducts to formation reaction product and reverse reaction - decomposing reaction product on mother substances. k 1 A + B + C = A 1 + B 1 + C 1 k 2

38. Classification of the chemical reaction according to the quantity of the reacting phases Homogeneous: N 2 (g) + H 2 (g) → NH 3 (g) Heterogeneous: Mg (s) + HCl (l) → MgCl 2 (l) + H 2 g) Topochemical ( in the hard phase)

39. 39 Classification of chemical reactions according to the molecularity Unimolecular (monomolecular): Н 2 СО 3 → Н 2 О + СО 2 Bimolecular: CuO + CO → Cu + CO 2 Termolecular: 2 NO + O 2 = 2 NO 2 The molecularity of an elementary reaction is the number of molecules coming together to react

40. The rate of a reaction is the speed at which a reaction happens.reaction

42. The Kinetic curves of the rate reaction’ mean value Initial substances Product of reaction Сonc entra tion t α Veritable (true) rates: dC = v·dt;v = tgα

43. FACTORS WHICH INFLUENCE RATES OF CHEMICAL REACTIONS 1. 1. Concentration of the reacting species. 2. 2. Temperature of the system. 3. 3. Nature of the reactants and products. 4. 4. Presence of a catalyst. 5. 5. Surface area. 6. 6. Exposure to radiation.

44. 1. 1. Concentration of the reactants. The rate of a reaction is directly proportional to the concentration of the reactants. Rate law expression

45. THE LAW OF MASSACTION (The Rate laws) The rate of reaction is proportional to the concentrations of reactants raised to а power. A + B = C;  = k[A][B] equation of the rate laws 3H 2 +N 2 =2NH 3 ;  = k[H 2 ] 3 [N 2 ] The coefficient k is called the rate constant for the reaction or velocity constant. The rate constant is independent of the concentrations but depends on the temperature. [A] = [B] = 1 mole/liter, then rate = k

46. 3. Temperature of the system. In general, an increase in the temperature increases the rate of almost all chemical reactions. This effect is observed for exothermic as well as for endothermic reactions. A general approximate rule for the effect of temperature on reaction rates is that the reaction rate for most of the chemical reactions becomes almost double, for every 10 0 C rise in temperature. This is also called temperature coefficient. It is the ratio of rate constants of the reaction at two temperature differing by 10 0 C Thus

47. or Vant-Hoff’s rule : where T 2  T 1 Temperature coefficient of reaction:

48. Arrhenius Equation It is a well-known fact that raising the temperature increases the reaction rate. E a = activation energy R = 8.314 [ J · mol -1 · K -1 ] T = absolute temperature in degrees Kelvin A = pre-exponential or frequency factor A = p · Z, where Z is the collision rate and p is a steric factor. Z turns out to be only weakly dependant on temperature. Thus the frequency factor is a constant, specific for each reaction.

49. 4. Presence of a catalyst. A catalyst is a substance which influences the rate of a reaction without undergoing any chemical change itself. It has been observed that many reactions are made to proceed at an increased rate by the presence of certain catalysts. 5. Surface area. The large the surface area of the reactants, the faster is rate of reaction. It has been observed that if one the reactants is a solid, then the rate of the reaction depends upon the state of sub-division of the solid. 6. Exposure to radiation. In some cases, the rate is considerably increased by the use of certain radiations. For example, reaction of hydrogen and chloride takes place very slowly in the absence of light. However, in the presence of light, the reaction takes place very rapidly.

50. Catalysis. A substance which changes the speed of a reaction without being used itself is called a catalyst. The phenomenon of increasing the rate of reaction by the use of catalyst is called catalysis. If а catalyst increases (accelerates) the speed of а reaction, it is called а positive catalyst and the phenomenon is called positive catalysis. On the other hand, if а catalyst decreases (retards) the speed of а reaction, it is called а negative catalyst and the phenomenon is called negative catalysis.

51. 1. Homogeneous catalysts. If the catalyst is present in the same phase as the reactants, it is called а homogeneous catalyst and this type of catalysis is called homogeneous catalysis. NO(g) 2 SO 2 (g) + О 2 (g) ===== SO 3 (g) Н + (aq) CH 3 COOC 2 H 5 (l)+Н 2 О(l)=====СН З СООН(l)+C 2 H 5 OH(1) Н + (aq) С 12 Н 22 О 11 (aq)+Н 2 О(1)====С 6 Н 12 О 6 (aq)+С 6 Н 12 О 6 (aq) Sucrose Glucose Fructose

52. 2. Heterogeneous catalysts. If the catalyst is present in а different phase than the reactants, it is called а heterogeneous catalyst and this type of catalysis is called heterogeneous catalysis. Pt, 800 0 С 4NH 3 + 5O 2 ======== 4NO + 6Н 2 O

53. Types of catalysis PositiveNegativeAutocatalysis Homogeneous HeterogeneousEnzyme Acid-base specificAcid-base unspecific

54. Enzymes Substance that acts as a catalyst in living organisms, regulating the rate at which life's chemical reactions proceed without being altered in the process. Enzymes are classified by the type of reaction they catalyze: 1. 1. Oxidation-reduction 2. 2. Transfer of a chemical group 3. 3. Hydrolysis 4. 4. Removal or addition of a chemical group 5. 5. Isomerization 6. 6. Polymerization

55. Influence on the activity of enzymes: 1. Enzyme activity can be affected by other molecules. Inhibitors are molecules that decrease enzyme activity; If a competing molecule blocks the active site or changes its shape, the enzyme's activity is inhibited. If the enzyme's configuration is destroyed (denaturated), its activity is lost. Activators are molecules that increase activity. Many drugs and poisons are enzyme inhibitors. 2. Activity is also affected by temperature 3. Chemical environment (pH). 4. The concentration of substrate.

56. The optimal meaning of рН for enzymes EnzymeSubstrateрН  -fructofuranozydaza Urease Papain Pepsin Arginase Saccharose Urea Protein Arginine 4,5-6,6 6,7 5,0 1,5-2,0 9,5-9,9