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Lecture 1. Chemical Thermodynamics and Bioenergetics. Fundamentals of Thermochemistry

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Lecture 1. Chemical Thermodynamics and Bioenergetics. Fundamentals of Thermochemistry Lecture 1. Chemical thermodynamics and bioenergetics. Fundamentals of thermochemistry. Lecture plan . 1. Types of thermodynamic systems. 2. Thermodynamics functions and parameters of system. 3. The first law of thermodynamics. Internal energy. Enthalpy. 4. Heat of isobaric and isochoric process. Standard heats of the substance formation and combustion. 5. Thermochemistry. Hess law. Thermochemical transformation. 6. Thermochemical calculation and their use for energetic characteristics of biochemical process. 7. Second law of thermodynamic. Entropy. Subject of thermodynamics All chemical reactions are accompanied by transformation of chemical energy to other forms of energy - thermal, electrical, mechanical, etc. Thermodynamics is the branch of physical science that studies all forms of energy and their mutual transformations; therefore it is sometimes called energetics. Bioenergetics is a field of thermodynamics that deals with biosystems. Classical thermodynamics is based on propositions which are confirmed by experiment and does not use knowledge about the molecular structure of substances.The energy of reactions is studied by the branch of thermodynamics which is called thermochemistry or chemical thermodynamics. In thermochemistry two types of chemical reactions are distinguished: exothermic (are accompanied by heat release) and endothermic (are accompanied by heat absorption). There are reactions (not so numerous), which are not accompanied by heat exchange. Chemical reactions can occur at a constant pressure (for example in an open flask) - these are isobaric processes, at a constant volume (in a closed flask or an autoclave) - these are isochoric processes, or at a constant temperature - these are isothermal processes (the names are derived from the Greek words isos - identical, baros - pressure, chorus - space, thermos - heat). Thermodynamics deals with the study of properties of various thermodynamic systems and processes occurring in them. A thermodynamic system is anybody or totality of bodies being in interaction with each other, which may be separated (conditionally or practically) from the surroundings for studying by thermodynamic methods. Different types of thermodynamic systems are known: 1. Homogeneous system – it is uniform in all its parts. For example an aqueous solution of ethanol, or a mixture of gases. 2. Heterogeneous system – it is not uniform and consists of two or more phases, e.g. water-benzene. The term phase means a part of a system with a characteristic chemical composition and macroscopic properties. Phases are separated from each other by physical surfaces, and at transition of these surfaces the properties sharply vary. 3. Physical system – it is a system, in which processes are accompanied by energy change, but the chemical nature of a substance is invariable. For example, changing of the modular condition of a substance at its melting (or crystallization) temperature; condensation of liquid vapor at its boiling temperature (water boils at 373 K). 4. Chemical system – it is a system, in which both phenomena take place: change of the energy content, change of the chemical nature of the system components. For example, interaction of zinc with sulfuric acid and other chemical reactions. 5. Open system – it is a system that may exchange both energy and substance with the surroundings. For example, a bio system - a living organism. 6. Closed system – it is a system that may exchange only energy with other systems, but not substance. For example, an electric range. 7. Isolated system – it is a system, which doesn’t exchange energy or substance with the surroundings. It is very difficult to create an absolutely isolated system. Reactors with good thermoisolation may be reckoned among these systems. At each moment of time the state of a system is characterized by physical properties that do not depend upon the previous history of the system state, for ex- ample: temperature T, pressure P, volume V, energy E, mass m, internal energy U, enthalpy H, entropy S, Gibbs energy G, Helmholtz energy F, etc. Thermodynamics is of great importance for medicine since it helps: • To generate scientific representation of the energy balance of a living organism. • To establish connection between the caloric content of food and energy expenses of the organism. • To develop objective criteria for determining the possibility of realization of separate processes in the human body without carrying out tests. There are some formulations of the first law of thermodynamics. Such concepts as «heat and work», on the one hand, and «internal energy and enthalpy», on the other hand, underlie it. Heat and work are different forms of energy transmission. In thermodynamics heat and work are algebraic values that may be positive and negative. Work is measured in joules. Heat is also expressed in joules in the SI, the unit calorie is also applied. The connection between a joule and a calorie is: 1.00 cal = 4.184 J. When heat is absorbed by a system from the surroundings, it has a positive value, if heat is released by a system into the surroundings, it is taken as negative. Using the concepts heat and work the first law of thermodynamics may be for- mulated as: Energy can neither be created nor destroyed, but only can be converted from one form into another (including heat and work), without changing quantitatively. In fact it is the law of conservation of energy, which was formulated by M. Lo- monosov as long ago as 1748. Other formulation of the first law is: It is impossible to develop a perpetuum mobile of the first kind (i.e. a machine producing work without expenditure of energy). It is possible to formulate the first law of thermodynamics on the basis of other reasons, introducing the concepts internal energy and enthalpy. Internal energy may be considered a sum of different types of energy from atoms, ions and molecules (energy of molecular motion, of intermolecular interac- tion, etc.). According to the law of energy conservation, the heat that is absorbed by a sys- tem is spent to change its internal energy and to produce work: Q = Δ U + A (1) For chemical processes the work against external forces is work against exter- nal pressure and it is equal: A = p(V2 - V1) = p ΔV (2) For an isochoric process (V— const): A = 0 and Qv = U2 - U1 = Δ U (3) It means that the system doesn’t produce external work that is associated with a volume change, and all heat that is released or absorbed is spent on changing the internal energy of a system. For an isobaric process (p — const), excluding internal energy changes, certain work (A) is carried out as a result of volume change in a system, which is equal to the product between pressure (p) and change of the system’s volume (V ): A= p ΔV (4) Qp = Δ U + p ΔV (5) Qp = (U2 - U1) + p(V2 - Vl) (6) or Qp = (U2 + pV2) - (U1 + pV1) (7) Assuming that U + pV = H (8) the heat of the processes taking place at constant temperature and pressure (the most widespread chemical processes) may be represented as: Qp = H2 - H1 = Δ H, (9) where H is the enthalpy of a system. The positive value of enthalpy change (ΔH > 0) corresponds to enthalpy in- crease or to heat absorption by a system (an endothermic process). The negative value of enthalpy change (ΔH< 0) corresponds to enthalpy decrease or to heat release by a system (an exothermic process). So in an isochoricprocess the heat of a reaction is equal to external energy change ΔU: QV = ΔU (10) and in an isobaric process heat is equal to a change of system’s enthalpy: Qp = ΔH (11) It must be noted that Qp > Qv on pΔV value which is the work of expansion. As well as the internal energy of a system U, enthalpy H is also a state function of a system. U and H may be considered as a measure of heat transportation at certain conditions: U at V = const, H at p = const. There is a relationship between internal energy and enthalpy of a system: ΔH = ΔU+pΔV From the equation (5) it follows: ΔU = Qp - pΔV (12) The equation (12) may be interpreted as a mathematical expression of the first law of thermodynamics. An increase of the internal energy of a system is equal to the heat, which is received by the system from the outside, except for the work produced by the system against external forces. All thermochemical calculations are based on Lavoisier and Laplace’s law and Hess’s law. The law of Lavoisier and Laplace (1780): The heat of decomposition of a chemical compound into simple substances is numerically equal and opposite in sign to the heat of formation of this compound from simple substances. Hess’s law (1840): The heat of a reaction is independent of the way, in which this reaction occurs, and only depends upon the initial and final states of a system. The consequences of Hess’s law are of great importance for thermochemical calculations. Consequences of Hess’s law: 1. Enthalpy of a forward reaction is equal and opposite in sign to enthalpy of a reverse reaction. Δ =Δ 2. Reaction enthalpy is equal to the sum of enthalpies of reaction products formation minus the sum of enthalpies of reactant formation. 3. Enthalpy of a combustion reaction is equal to the sum of enthalpies of reactant combustion (Δ ) minus the sum of enthalpies of product combustion. ΔHc = - For a reaction nA + mB = qC + pD Δ Hf = [qΔH°f C + pΔH°f D] - [nΔH°f A + mΔH°f B]; Δ Hc = [nΔHc A + mΔHc B] - [qΔHc C + pΔHc D]. The data on the thermal effects of reactions are used for: calculation of the thermal balances of technological processes, determination of the energy of inter- atomic and intermolecular bonds, ascertainment of the structure and reactionary ability, establishment of the direction of chemical processes, description of the energy balance of an organism.
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