Lecture 1. Chemical Thermodynamics and Bioenergetics. Fundamentals of Thermochemistry

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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 example: 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 formulated 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 interaction, etc.). According to the law of energy conservation, the heat that is absorbed by a system 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 external 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 increase 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.

Second law of thermodynamics

All the processes connected with transition of one type of energy to another refer to the first law of thermodynamics, the law of energy conservation. However, it is important to know not only the energy of processes (for example, the heat of formation or substance decomposition), but also which factors influence the direction and depth of the proceeding of chemical reactions. Another no less important question is whether a given reaction will occur spontaneously, without external intervention, or not. The answers to these questions are given by the second law of thermodynamics.

There are some formulations of the second law:

It is impossible to construct a perpetuum mobile of the second kind, i.e. it is impossible to transform heat into work completely. ( W. Thomson )

The work of each electric power station leads to thermal contamination of the surroundings because some part of energy is being lost.

It is impossible to transfer heat from a cooler body to a hotter body without performing work. (E.R. Clausius)

The refrigerator’s temperature will stay less than the outdoor temperature only if electric energy is spent.

A process, which under particular conditions occurs by itself without an extraneous source of energy, is called spontaneous. For example: falling of a stone from hands to the floor, expansion of an ideal gas, melting of ice, dissolution of salts, evaporation of liquids, etc.

In the adduced examples the motive power of the processes consists in transition from a thermodynamic system with a more regulative state into a less regulative state. Reverse transition of a system is hardly probable.

For a quantitative estimation of the probability of a system state or for an estimation of the disorder degree a thermodynamic function as entropy S has been proposed. Entropy is a measure of a system disorder. Entropy is a state function: its change (ΔS) depends only on the initial and final states of a system.

Entropy is connected with the thermodynamic probability of the realization of some particular system state by L. Boltzmann’s equation:

S = KlnW,

where K is Boltzmann’s constant, W is the thermodynamic probability or the number of possible microstates which may be realized for a particular system macrostate.

Entropy is measured in J/mol • K.

From Boltzmann’s law it follows that the entropy of a pure ideal crystal is equal to zero at the temperature of absolute zero (W = 1 for it, then S = K In 1 = 0). It is the most regulated system. In other systems the value W is greater and S > 0. The bigger the system disorder, the higher its entropy. Entropy is connected with the thermal characteristic of a system by the following correlation:

The product TΔS is called connected energy.

The concept of entropy underlies the second law of thermodynamics:

In isolated systems, processes occur spontaneously on condition of entropy increase. In the real world isolated systems are found very rarely. In real systems processes may be accompanied by both an increase and a decrease of entropy.

An important conclusion follows from the second law of thermodynamics:

The total change in entropy that is necessary for the formation of a human body and maintenance of its life and the life of any other living system is always positive.

The human body is a complex, highly organized, and very regulated system. Its entropy is much less than the entropy of the same quantity of CO2, H2O and some other substances, which compose the organism. But proceeding of many thousands of chemical reactions that are necessary for the recreation and vital activity of the organism are accompained by a substantial increase of entropy in the environment.

All human activity aimed at ruling the world that surrounds us demands high energy expenditure, which eventually increases disorder.

At various transformations it is important to know not the absolute value of entropy, but its change ΔS. As well as enthalpy change (ΔH), entropy change (ΔS!) may be calculated using the following equation:

0 0 0 ΔS = Σ S prod – Σ S react

To compare the entropy of different substances, these, as well as enthalpy of formation, are given under standard conditions (S°):

T = 298 K, p = 101.3 kPa, n(x) = 1 mol.

The value of entropy allows:

1) to forecast which processes may occur spontaneously and which cannot;

2) to predict the direction of possible transformations and to control them. The dependence of entropy on temperature is formulated in the third law of thermodynamics, or Nernst’s heat theorem, or Planck’s postulate:

The entropy of a pure ideal crystal at absolute zero is equal to zero.

Application of II law of thermodynamics to biological systems

A. Although living organisms is open, non equilibrium systems are applicable to them, and I and II of the laws of thermodynamics, as biochemical processes are irreversible, occur spontaneously. In other words living organisms is stationary systems. Part of the energy that is released during the oxidation of food, irreversibly converted into heat, which dissipates into the surrounding space.

B. In the body, all the processes are spontaneous and therefore the entropy S increases. But the body temperature does not rise and does not come "heat death" because body consumes a substance with low entropy (IUD), and highlights the decay products with high entropy (small molecules). As a result, the entropy of an open system – is a constant value.