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LECTURE № 6 THEME: Theoretic bases of bioenergetics. Electrochemistry. associate. prof. Yevheniy. B. Dmukhalska 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. 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: 3 Common terms of thermodynamics System Parameter Condition (state) Process 4 System is a specified part of the universe which is under observation 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 KNO3 heterogeneous KNO3 PbI2↓ Types of Systems 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. 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 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 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 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 14 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) 15 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. 16 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. 17 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. 18 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. 19 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. 20 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 21 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. The meaning of the state functions in the thermodynamic processes Exothermic process Qv > 0, ▲U < 0 Qp > 0, ▲H < 0 Endothermic process Qv < 0, ▲U > 0 Qp < 0, ▲H > 0 22 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 23 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 24 Thermochemistry The study of the energy transferred as heat during the course of chemical reactions. Thermochemical reactions: H2(g) + Cl2(g) = 2HCl; ▲ H = -184,6 kJ 1/2 H2(g) + 1/2 Cl2(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 25 The Hess’s law Initial reactants Н2 Н1 Н3 Н4 The products of reaction Н1 = Н2 + Н3 + Н4 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. 26 Conclusions from the Hess law 1.Нc298(the standard enthalpy of combustion) =Нf298(the standard enthalpy of formation) 2.Н(formation)= ΣnНf298(products) - ΣnНf298(reactants) 3.Н(combustion) = ΣnНс298(reactants) - ΣnНс298(products) 4.For elementary substances Н0298 = 0 4.Н3=Н1-Н2 5.Н1=Н3-Н2 1 1 2 3 2 3 27 Correlation U і Н: H U pV U RT If υ0, so НU: СаО + СО2 → СаСО3 If υ0, so НU: Na + H2O → NaOH + H2 If υ=0, so Н=U: H2 + Cl2 → 2HCl 28 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. 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 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 31 1) Spontaneous (irreversible) process : ▲ G < 0, ▲S > 0, ▲H < 0 2) Unspontaneous (reversible) process : ▲ G > 0, ▲S < 0, ▲H > 0 3) Equilibrium state ▲G=0 32 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. Temperature approaches absolute zero (0 K), the entropy of a system approaches a constant (and minimum) value. The entropy of a perfect crystalline state is zero at 0 K. 33 Electrochemistry The branch of science, which deals with the study oxidation-reduction reaction to produce the interconversion of chemical and electricl energy. of transition chemical energy to electrical energy is known as electrochemistry. 34 Nernst’s equation Мn++nе = М Then the Nernst eqn. is applied as follows: E = E0 – (RT/ nF) ln ([M]/ [Mn+]) where Е = electrode potential under given concentration of Мn+ ions and temperature Т; Е0 – standard electrode potential; R – gas constant, R = 8.315 J/K .mol; Т – temperature in К; F – Faradays constant, F = 96,485 С /mol; n – number of electrons involved in the electrode reaction. 35 Standard (normal) hydrogen electrode Pt, Н2 (g)/Н+ (Concentration) H2 = 2H+ + 2е2H+ + 2е- = H2 E = E0 – (RT/ 2F) ln (pH2/ [H+]2), E0H+/H2 = 0V. In the standard hydrogen gas electrode, hydrogen at atmospheric pressure is passed into 1 М НС1 in which foil of the platinized platinum remains immersed through which inflow or outflow of electrons takes place. 36 Since а cathode reaction is а reduction, the potential produced at such an electrode is called а reduction potential. Similarly, the potential produced at an anode is called an oxidation potential. These are known as standard reduction potentials or standard electrode potentials. They are usually tabulated for 25 С. 37 Types of electrodes 1. Metal-metal ion electrodes 2. Gas-ion electrodes 3. Metal-insoluble salt-anion electrodes 4. Inert "oxidation-reduction" electrodes 5. Membrane electrodes 39 Electrochemical cells allow measurement and control of a redox reaction. Electrodes of the first kind. An electrode of the first kind is а piece of pure metal that is in direct equilibrium with the cation of the metal. А single reaction is involved. For example, the equilibrium between а metal Х and its cation Х+n is: Х+n + ne- = X (s) for which Еnd = Е0X+n 0.0592 1 0.0592 – -------- log ---- = Е0X+n + ---------- log aX+n n aX+n n 41 The metal - metal ion electrode consists of а metal in contact with its ions in solution. An example: silver metal immersed in а solution of silver nitrate As a cathode: the diagram: Ag+(aq) Ag(s) half-reaction equation is: Ag+ (aq) + e-Ag(s) as an anode: the diagram: Ag(s) Ag+(aq) half-reaction equation is: Ag(s) Ag+(aq) + еNernst’s equation: E = E0 – (RT/ nF) ln ([Ag]/ [Agn+]) 43 Electrodes of the Second type. Metals not only serve as indicator electrodes for their own cations but also respond to the concentration of anions that form sparingly soluble precipitates or stable complexes with such cations. AgCl + e- = Ag (s) + Cl E0AgCl = 0.222 V The Nernst expression for this process is: EAgCl = E0AgCl – 0.0592 log [Cl-] = 0.222 + 0.0592 pCl 44 In the metal-insoluble salt-anion electrode, а metal is in contact with one of its insoluble salts and also with а solution containing the anion of the salt. An example is the so-called silver - silver chloride electrode, written as а cathode as: Cl- (aq) AgCl(s) Ag(s) for which the cathode half-reaction is: AgCl (s) + е- Ag(s) + Cl- (aq) Nernst’s equation: E = E0 – (RT/ 1F) ln ([Ag] [Cl-]/ [AgCl]) 46 An inert oxidation-reduction electrode consists of а strip, wire, or rod of an inert materiel, say, platinum, in contact with а solution, which contains ions of а substance is two different oxidation states. Thus, for the ferric - ferrous ion electrode functioning as а cathode, Fe3+, Fe 2+(aq) Pt(s) the iron(III), or ferric, ion, Fe+3(aq), is reduced to the iron(II), or ferrous, ion, Fe+2(aq): Fe+3(aq) + е- Fe+2(aq) Nernst’s equation: E = E0 – (RT/ 1F) ln ([Fe+2]/ [Fe+3]) 47 а membrane electrode - the glass electrode. This can be depicted as: Pt(s) Ag(s) AgC1(s) HC1(aq,1M) glass Cell can be depicted as: reference electrode salt bridge analyte solution indicator electrode Ecell = Eind + Eref + Ej 49 1.Glass electrode – indicator electrode; diagram which is: Ag(s) AgC1(s) HC1(aq,1M) glass 2.Bulb of glass electrode. 3.Solution of unknown pH. 4.Silver-silver chloride electrode electrode;diagram which is: Cl- (aq) AgCl(s) Ag(s) 5.Amplifying potentiometer. reference 50 Cell potential or EMF of a cell. The difference between the electrode potentials of the two half cell is known as electromotive force (EMF) of the cell or cell potential or cell voltage. The EMF of the cell depends on the nature of the reactants, concentration of the solution in the two half cells, and temperature. 53 Reference electrode is electrode potential which stabile А hydrogen electrode is seldom used as а reference electrode for day-today potentiometric measurements because it is somewhat inconvenient and is also а fire hazard. 54 An ideal indicator electrode responds rapidly and reproducibly to changes in the concentration of an analyte ion (or group of ions). Although no indicator electrode is absolutely specific in its response, а few are now available that are remarkably selective. There are two types of indicator electrodes: metallic and membrane. Metallic indicator electrodes: Electrodes of the first kind. Electrodes of the Second Kind. Membrane Electrodes 55 The relationship between pH and the voltage of the hydrogen elect calomel electrode cell at 250С can be written as Ecell Ecalomel 1 pH = --------- - (---------- + ---- log pH2) 0.0592 0.0592 2 Ecell pH = ---------- = constant 0.0592 56