Direction of spontaneous flow of processes in an isolated system. Levchenkov S.I. Catalysis is homogeneous and heterogeneous. Enzyme catalysis. Michaelis-Menten equation

Section II. solutions and heterogeneous equilibria

Basic concepts and definitions

Substances that form a thermodynamic system can be in different states of aggregation: gaseous, liquid, solid.

A thermodynamic system within which there are no interfaces separating parts of the system that are different either in physical structure or in chemical properties is called homogeneous.

A thermodynamic system consisting of parts with different physical or chemical properties, separated from each other by interfaces, is called heterogeneous.

Any heterogeneous system consists of several phases.

Phase- this is part of a heterogeneous system, limited by the interface and characterized by the same physical and chemical properties at all points.

There are single-phase, two-phase, three-phase, etc. systems.

Every system consists of one or more substances called components.

Components– individual substances that make up the system, and which can be isolated from the system and exist outside of it.

Number independent components is the smallest number of individual substances required to form a given system. It is equal to the total number of individual substances included in a given system, minus the number of equations connecting these substances.

Based on the number of components, one-component, two-component, three-component, etc. are distinguished. systems.

Any system is characterized by external and internal state parameters.

The number of independent thermodynamic parameters of a given system, the derivative change of which, within certain limits, does not cause the disappearance of some and the formation of other phases is called number of thermodynamic degrees of freedom, or variability, systems.

Based on the number of thermodynamic degrees of freedom, systems are divided into invariant ( WITH= 0), monovariant ( WITH= 1), divariant ( WITH= 2), etc.

Solution is a homogeneous single-phase system consisting of at least two independent components, each elementary volume of which has the same physical, chemical and thermodynamic properties.

Solvent Usually a substance is considered, the amount of which in a solution is greater or which does not change its state of aggregation during the formation of a solution, the remaining components are called dissolved.

There are solid, liquid and gaseous perfect And real solutions.

Ideal A solution is called such that all components are characterized by the same shape and size of molecules and the same energy of intermolecular interactions.

Ideal solutions are quite rare. These are homogeneous mixtures of substances with similar physical and chemical properties. For example, mixtures of optical isomers, neighboring members of the same homologous series. The model of an ideal gas solution is a mixture of ideal gases. Ideal solutions often include infinitely dilute solutions.

Most solutions are real.

Real solutions are solutions whose components differ either in shape or size or in the energy of intermolecular interactions.

All properties of solutions are divided into extensive And intensive.

Extensive properties - properties that depend both on the total mass of the solution and on its composition, for example V, U, H, G, S, C p.

These properties relate to the entire solution as a whole, and not to its individual components.

Intensive properties are properties that depend only on the composition of the solution and do not depend on its total mass, for example, saturated vapor pressure.

To characterize solutions use average moles And partial moles properties.

Average molar property– extensive property of 1 mole of solution.

For example, the average molar volume can be calculated using the formula:

Where n 1 , n 2 , n 3,... – the number of moles of the first, second, third, etc. components.

Partial molar property th component is the partial derivative of the extensive property of the solution with respect to the number of moles of this component ( n i) with a constant quantity of all other components and external parameters ( R And T).

The partial molar property is a characteristic of an individual component of the system, i.e. is an intensive property of a solution.

For example, the partial molar volume i th component is the partial derivative

.

Gibbs free energy(or simply Gibbs energy, or Gibbs potential, or thermodynamic potential in the narrow sense) is a quantity that shows the change in energy during a chemical reaction and thus gives an answer to the question about the fundamental possibility of a chemical reaction occurring.

direction of a chemical reaction defines Gibbs energy(∆G). Gibbs energy is also called isobaric-isothermal potential. The dimension of the Gibbs energy is kJ/mol.

At constant pressure and temperature ( р=const, T=cons)t the reaction proceeds spontaneously in the direction corresponding to the decrease in Gibbs energy. If ∆G< 0, then the reaction spontaneously proceeds in the forward direction. If ∆G > 0 , then the spontaneous occurrence of the process in the forward direction under these conditions is impossible, but the reverse process is possible. If ∆G = 0 , then the reaction can proceed both in the forward direction and in the reverse direction, and the system is in a state of equilibrium.

The change in Gibbs energy during a chemical reaction (∆) does not depend on the path of the process and can be calculated by a consequence of Hess’s law: Gibbs energy change as a result of a chemical reaction is equal to the sum of the Gibbs energies of the reaction products minus the sum of the Gibbs energies of the starting substances, taking into account the stoichiometric coefficients. For example, the standard Gibbs energy of a reaction

aA + bB = cC + dD

where ∆G 0 is the standard Gibbs energy of formation of a substance, kJ/mol.

The Gibbs energy of formation of simple substances is zero. ∆ has the same dimension as enthalpy and is therefore usually expressed in kJ.

The change in the standard Gibbs energy of a chemical reaction can also be calculated using the equation:

∆ = ∆ – Т∆ , Where

T – absolute temperature,

∆ – change in entropy.

∆H c.r. – enthalpy change.

During a chemical interaction, enthalpy, which characterizes the heat content of the system, and entropy, which characterizes the system’s tendency toward disorder, simultaneously changes. A decrease in enthalpy and an increase in entropy are two driving forces of any chemical process. In a state of balance ∆ =0, Means:

∆ – Т∆ =0 And

If we neglect changes in ∆H 0 h.r. and ∆S 0 x.r with increasing temperature, then we can determine the temperature at which equilibrium of the chemical reaction is established for the standard state of the reagents:

T equal =

Many chemical reactions occur spontaneously, i.e. without external energy consumption. One of the driving forces of a spontaneous chemical process is a decrease in the enthalpy of the system, i.e. exothermic heat effect of a reaction. The other is the tendency of particles (molecules, ions, atoms) to chaotic movement, disorder. A measure of the chaotic, disordered state of a system is a thermodynamic function called entropy (S).

When a system transitions from a more ordered state to a less ordered state (heating, evaporation, melting), entropy increases (DS>0). In the case of a system transition from a less ordered state to a more ordered one (cooling, condensation, crystallization), the entropy of the system decreases (DS<0).

In isolated systems, only those processes occur spontaneously that are accompanied by an increase in entropy (S>0)- This is the essence of the second law of thermodynamics.

The entropy of a substance in the standard state is called standard entropy (So) and has the unit J/mol K

The entropy of a substance in the gaseous state is significantly higher than in the liquid and solid states, therefore the change in entropy in a chemical reaction is judged by the change in the number of moles of gaseous substances.

The possibility of spontaneous occurrence of a chemical process is determined by two factors:

The desire to form strong bonds between particles, to the emergence of more complex substances, which is accompanied by a decrease in the energy of the system - enthalpy factor (DH<0);

The desire for separation of particles, for disorder, which is characterized by an increase in entropy - entropy factor (DS>0).

These factors are united by a function called Gibbs energy(DG), equal to: DG = DH - T DS. (D is delta type, triangle is shorter)

The change in the Gibbs energy serves as a criterion for the spontaneous occurrence of a chemical reaction:

A chemical reaction is fundamentally possible if the Gibbs energy decreases during the reaction (DG<0);

A chemical reaction cannot proceed spontaneously; if the Gibbs energy of the system increases (DG>0), a reverse reaction occurs;

A chemical reaction can occur in both forward and reverse directions, i.e. the system is in a state of equilibrium (DG=0).

From the equation DG=DH-T DS it follows:

If DН<0 и DS>0, then always DG<0, т.е. реакция с выделением теплоты и увеличением степени беспорядка возможна при любых температурах;

If DH>0 and DS<0, то всегда DG>0, i.e. a reaction with the absorption of heat and an increase in the degree of order is impossible under any conditions;

DH>0,DS<0. Реакция будет протекать в прямом направлении только при условии, что |T DS|>|DH|. These reactions occur at high temperatures;

D.H.<0, DS>0. Condition for spontaneous reaction: |DH|>|T DS|. Such reactions usually occur at low temperatures.

The temperature at which the sign of the Gibbs energy of a reaction changes can be determined from the equilibrium condition:

Tr = DH/DS, where Tr is the temperature at which equilibrium is established.

The change in the Gibbs energy of a system when 1 mole of a substance is formed from simple substances that are stable under standard conditions is called the standard Gibbs energy of formation of a substance (DGof). The standard Gibbs energy of formation of simple substances is taken to be zero.

The standard Gibbs energy of a chemical reaction (DGor) can be calculated as the sum of the standard Gibbs energies of formation of reaction products minus the sum of the Gibbs energies of formation of starting substances, taking into account stoichiometric coefficients.

Second law of thermodynamics: under isobaric-isothermal conditions (p, T = const), only such processes can spontaneously occur in the system, as a result of which the Gibbs energy of the system decreases (ΔG< 0). В состоянии равновесия G = const, G = 0

A reversible process – if, during the transition from the initial state to the final state, all intermediate states turn out to be in equilibrium

An irreversible process – if at least one of the intermediate states is nonequilibrium.

Entropy– a state function, the increment of which ΔS is equal to the heat Qmin supplied to the system in a reversible isothermal process, divided by the absolute temperature T, three of which the process is carried out: ΔS = Qmin/ T or a measure of the probability of the system being in a given state - a measure of the disorder of the system.

Gibbs energy– state function, which is a criterion for the spontaneity of processes in open and closed systems.

ΔG = ΔH – TΔS

ΔrG = ΔrG0 + RTln ca(A)cb(B) – van’t Hoff isotherm

The chemical potential of a substance X in a given system is a value determined by the Gibbs energy G (X) per mole of this substance. μ(X) = n(X)

The criteria for the direction of spontaneous occurrence of irreversible processes are the inequalities ΔG< 0 (для закрытых систем), ΔS >0 (for isolated systems).

During a spontaneous process in closed systems, G decreases to a certain value, taking the minimum possible value for a given system Gmin. The system goes into a state of chemical equilibrium (ΔG = 0). The spontaneous course of reactions in closed systems is controlled by both enthalpy (ΔrH) and entropy (TΔrS) factors. For reactions for which ΔrH< 0 и ΔrS >0, the Gibbs energy will always decrease, i.e. ΔrG< 0, и такие реакции могут протекать самопроизвольно при любых температурах

In isolated systems, entropy is the maximum possible value for a given system, Smax; in equilibrium ΔS = 0

Standard Gibbs energy: ΔrG = ΣυjΔjG0j – ΣυiΔiG0i

Second law (second law) of thermodynamics. Entropy. The release of thermal energy during a reaction helps it proceed spontaneously, i.e. without outside interference. However, there are other spontaneous processes in which heat is zero (for example, the expansion of a gas into a void) or even absorbed (for example, the dissolution of ammonium nitrate in water). This means that in addition to energy factor the possibility of processes occurring is influenced by some other factor.

It is called entropy factor or entropy change. Entropy S is a function of state and is determined by the degree of disorder in the system. Experience, including everyday experience, shows that disorder arises spontaneously, and that in order to bring something into an orderly state, you need to expend energy. This statement is one of the formulationssecond law of thermodynamics.

There are other formulations, for example: Heat cannot spontaneously transfer from a less heated body to a more heated one (Clausius, 1850). A block heated at one end eventually takes on the same temperature along its entire length. However, the reverse process is never observed - a uniformly heated block does not spontaneously become warmer at one end and colder at the other. In other words, the thermal conduction process is irreversible. To take away heat from a colder body, you need to expend energy; for example, a household refrigerator uses electrical energy for this.

Consider a vessel divided by a partition into two parts filled with different gases. If you remove the partition, the gases will mix and never separate spontaneously again. Add a drop of ink to a container of water. The ink will be distributed throughout the entire volume of water and will never spontaneously collect into one drop. In both cases, spontaneous processes are accompanied by an increase in disorder, i.e. increase in entropy (Δ S> 0). If we consider an isolated system, the internal energy of which cannot change, then the spontaneity of the process in it is determined only by the change in entropy: In an isolated system, only processes accompanied by an increase in entropy occur spontaneously (Boltzmann, 1896). This is also one of the formulations of the second law of thermodynamics. The manifestation of the entropy factor can be clearly seen in phase transitions ice–water, water–steam, leaking at constant temperature. As is known, absorption (ice drift - cooling) and heat release (freezing - warming) occur in such a way that

Δ HГ = Δ S×T,

where Δ H G – “latent” heat of phase transition. In phase transitions ice–water–steam – Δ S l < ΔS V < ΔS p, i.e. entropy increases when moving from a solid to a liquid and from a liquid to a gas, and its value increases the more randomly the molecules move. Thus, entropy reflects the structural differences of the same chemical element, molecule, or substance. For example, for the same water, H2O is a crystal, liquid, steam; for carbon – graphite, diamond, allotropic modifications, etc.

The absolute value of entropy can be estimated using third law of thermodynamics(Planck's postulate), which states that the entropy of an ideal crystal at 0 K is equal to zero lim Δ S=0 (at T=0).

The SI unit of entropy is J/K×mol. It is clear that absolute zero temperature is unattainable (a consequence of the second law of thermodynamics), but it is important in determining temperature - the Kelvin scale.

In addition, there is another function of the state of matter - heat capacity

Δ C = Δ H/Δ T,

which has the same dimension as entropy, but means the ability of a substance to give (receive) heat when the temperature changes. For example, when the temperature changes by 100°C, steel will heat up and cool down faster, respectively, than brick, which is why stoves are made of brick. Specific or molar heat capacity values ​​are tabulated in reference books, usually for isobaric conditions - cp.

Standard absolute entropies S°298 formations of some substances are given in reference literature. Please note that unlike Δ Hf simple substances matter S°298 > 0, because their atoms and molecules are also in random thermal motion. To find the entropy change in a reaction, you can also apply a corollary of Hess's law:

Δ S(reactions) = Δ S(products) – Δ S(reagents)

Δ S> 0 according to the second law of thermodynamics favors the reaction, Δ S < 0 - препятствует.

It is possible to qualitatively estimate the sign of Δ S reactions based on the aggregate states of reactants and products. Δ S> 0 for the melting of solids and evaporation of liquids, dissolution of crystals, expansion of gases, chemical reactions leading to an increase in the number of molecules, especially molecules in the gaseous state. Δ S < 0 для сжатия и конденсации газов, затвердевания жидкостей, реакций, сопровождающихся уменьшением числа молекул.

Using reference data, we calculate Δ S°298 reaction (a).

S°298 4×28.32 3×76.6 2×50.99 3×64.89

Δ S°298 = 2×50.99+3×64.89–4×28.32–3×76.6 = –46.4 J/K.

Thus, the entropy factor prevents the occurrence of this reaction, and the energy factor (see above) is favorable.

Is there really a reaction? To answer this question, we need to simultaneously consider both factors: enthalpy and entropy.

Gibbs free energy. Criteria for the spontaneous occurrence of chemical reactions. Simultaneous consideration of energy and entropy factors leads to the concept of another complete state function - free energy. If measurements are carried out at constant pressure, then the quantity is called Gibbs free energy(in old chemical literature - isobaric-isothermal potential) and is denoted by Δ G.

Gibbs free energy is related to enthalpy and entropy by the relation:

Δ G = Δ H – TΔ S

Where T– temperature in Kelvin. The change in the Gibbs free energy during the reaction of the formation of 1 mole of a substance from simple substances in standard states is called free energy of formation Δ G° and is usually expressed in kJ/mol. The free energies of formation of simple substances are assumed to be zero. To find the change in the Gibbs free energy during a reaction, you need to subtract the sum of the free energies of formation of the reactants from the sum of the free energies of formation of the products, taking into account the stoichiometric coefficients:

Δ G(reactions) = SΔ G(products) – SΔ G(reagents)

Spontaneous reactions correspond to Δ G < 0. Если ΔG> 0, then the reaction under these conditions is impossible. Consider reaction (a)

4Al (solid) + 3PbO2 (solid) = 2Al2O3 (solid) + 3Pb (solid)

Δ G°298 0 3×(–219.0) 2×(–1576.5) 0

Δ G°298 = 2×(–1576.5)–3×(–219.0) = –2496 kJ.

There is another way to calculate Δ G reactions. Above we found the values ​​of Δ H and Δ S, then Δ G = Δ H – TΔ S

Δ G°298 = –2509.8 kJ – 298.15 K×(–0.0464 kJ/K) = –2496 kJ.

Thus, reaction (1) under standard conditions proceeds spontaneously. Sign Δ G shows the possibility of carrying out the reaction only under the conditions for which the calculations were carried out. For a deeper analysis, it is necessary to separately consider the energy and entropy factors. There are four possible cases:

Table 7.3.

Determining the possibility of a chemical reaction occurring

Criteria

Opportunity

Δ H < 0, ΔS > 0

Both factors favor the reaction. As a rule, such reactions occur quickly and completely.

Δ H < 0, ΔS < 0

The energy factor favors the reaction, the entropic factor hinders it. The reaction is possible at low temperatures.

Δ H > 0, Δ S > 0

The energy factor hinders the reaction, the entropic factor favors it. The reaction is possible at high temperatures.

Δ H > 0, Δ S < 0

Both factors interfere with the reaction. Such a reaction is impossible.

If under standard conditions Δ G reaction > 0, but the energy and entropy factors are directed in the opposite direction, then it is possible to calculate under what conditions the reaction will become possible. Δ H and Δ S Chemical reactions themselves are weakly dependent on temperature if any of the reactants or products do not undergo phase transitions. However, in addition to Δ, the entropy factor S also includes absolute temperature T. Thus, with increasing temperature, the role of the entropy factor increases, and at temperatures higher T » Δ H/Δ S the reaction begins to go in the opposite direction.

If Δ G= 0, then the system is in a state of thermodynamic equilibrium, i.e. Δ G– thermodynamic criterion for chemical equilibrium of reactions (see the above phase transitions of water).

So, by analyzing the functions of the state of the system - enthalpy, entropy and Gibbs free energy - and their changes during a chemical reaction, it is possible to determine whether this reaction will occur spontaneously.

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Size: 82 KbThe first law of thermodynamics. Enthalpy. Standard enthalpy of formation of a substance. Standard enthalpy of combustion of a substance. Standard enthalpy of reaction. Hess's law. Application of the first law of thermodynamics to biosystems. Second law of thermodynamics. Entropy. Gibbs energy. Predicting the direction of spontaneous processes in isolated and closed systems. Examples of exergonic and endergonic processes occurring in the body. The principle of energy coupling Classification of reactions used in kinetics: homogeneous, heterogeneous, microheterogeneous; simple and complex (parallel, serial, conjugate, chain) Dependence of reaction rate on concentration. Molecularity of an elementary reaction act. Order of reaction. Kinetic equations of first and zero order reactions. Half-life period. Dependence of reaction rate on temperature. Temperature coefficient of reaction rate and its features for biochemical processes. Activation energy. Catalysis is homogeneous and heterogeneous. Enzyme catalysis. Michaelis-Menten equation. Chemical balance. Reversible and irreversible reactions. Philosophical worldview of Nikolai Aleksandrovich Berdyaev The thinker himself characterized it as “a philosophy of the subject, a philosophy of spirit, a philosophy of freedom, a dualistic-pluralistic philosophy, a creative-dynamic philosophy, a personalistic philosophy and an eschatological philosophy.” Turret lathe A turret lathe is used for processing workpieces or parts made from calibrated rods. The economic role of the state in the market economy The evolution of theoretical manifestations on the role of power in the economy. State regulation of the economy: subjects, objects and functions. Methods of sovereign regulation of economic processes. Social policy of the state Family sociogram Purpose of application: to identify the subject’s position in the system of interpersonal relationships and determine the nature of communication in the family. Coase theorem. Proof of the theorem Russian privatization in the light of the Coase theorem. Transaction costs. Essence and purpose. Transaction cost theory: the role of information costs.

Enthalpy and entropy factors characterizing two opposite tendencies of processes - the desire for unification, order and the desire for separation, disorder, taken separately cannot be criteria for the spontaneous flow of chemical reactions. For isobaric-isothermal processes they are united by a function called change Gibbs energy during the process or isobaric-isothermal potential(ΔG), equal to:

This equation can be written as:

As you can see, the thermal effect of a chemical reaction includes two parts. The first part of ΔG is equal to the maximum work Wmax, which the system can accomplish when the process is carried out in equilibrium under isobaric-isothermal conditions. Therefore, the change in the Gibbs energy of a reaction is part of the energy effect of a chemical reaction that can be converted into work:

Since the change in Gibbs energy of a reaction can be converted into work, it is also called free energy. The second term on the right side of the equation (entropy factor) represents the part of the energy effect that can be converted into heat dissipated into the environment. Therefore, the entropy factor is called bound energy.

The change in the Gibbs energy serves as a criterion for the spontaneous occurrence of chemical reactions during isobaric-isothermal processes. A chemical reaction is fundamentally possible if the Gibbs energy of the system decreases, i.e.

ΔG< 0.

This equation is a condition for the possibility of spontaneous reaction in the forward direction. A chemical reaction cannot proceed spontaneously if the Gibbs energy of the system increases, i.e.

ΔG>0.

This equation serves as a thermodynamic condition for the possibility of spontaneous occurrence of the reverse reaction. If

then the system is in equilibrium, the reaction proceeds in both forward and reverse directions.

The direction of chemical reactions depends on their nature. So the condition ΔG<0соблюдается при любой температуре для экзотермических реакций (ΔH<0), у которых в ходу реакции возрастает число молей газообразных веществ, и.следовательно, энтропия (ΔS>0). In such reactions, both driving forces are directed towards the forward reaction and ΔG<0 при любых температурах. Такие реакции самопроизвольно могут идти только в прямом направлении, т.е. являются необратимыми.

On the contrary, an endothermic reaction (ΔH > 0), during which the number of moles of gaseous substances decreases (ΔS<0), не может протекать самопроизвольно в прямом направлении при любой температуре, т.к. всегда ΔG> 0.

The possibility of many reactions occurring depends on temperature, since temperature affects the sign of the change in the Gibbs energy of these reactions. If as a result of an exothermic reaction (ΔH<0) уменьшается число молей газообразных веществ и соответственно энтропия системы (ΔS<0), то при вы невысоких температурах |ΔH| >|ТΔS| and the reaction can proceed spontaneously in the forward direction (ΔG< 0). При высоких же температурах |ΔH|<|ТΔS| и прямая реакция протекать не может, а обратная возможна.

To determine the temperature above which the sign of the change in the Gibbs energy of a reaction changes, you can use the condition:

where T r is the temperature at which equilibrium is established, i.e. the possibility of direct and reverse reactions occurring is equally probable.

If, as a result of an endothermic reaction (ΔH> 0), the number of moles of gaseous substances increases (ΔS> 0), then at low temperatures, when |ΔH| > |TΔS|, the direct reaction cannot occur spontaneously (ΔG> 0), and at high temperatures (T>T p) the direct reaction can occur spontaneously (ΔG< 0).

Table 3. Effect of temperature on the direction of chemical reactions

ΔH	ΔS	ΔG	Direction of reaction	Example
ΔH<0	ΔS>0	ΔG<0	The direct reaction can occur spontaneously at any temperature	C + 1/2O 2 = CO
ΔH>0	ΔS<0	ΔG>0	A direct reaction cannot occur spontaneously at any temperature	CO = C + 1/2O 2
ΔH<0	ΔS<0	ΔG<0 при Т0 at T>T p	A direct reaction can occur spontaneously at low temperatures and a reverse reaction at high temperatures.	CaO + CO 2 = CaCO 3
ΔH>0	ΔS>0	ΔG>0 at T Tp	A direct reaction can occur spontaneously at high temperatures and a reverse reaction at low temperatures.	CH 4 + 2H 2 O(g) = CO 2 + 4H 2

The Gibbs energy is a function of state and does not depend on the method of carrying out the process, but is determined only by the initial and final state of the system. The change in the Gibbs energy of a reaction obeys Hess's law and its corollaries, so it can be calculated using the equation:

The Gibbs energy of formation of simple substances is zero. If a substance is in a standard state, then the Gibbs energy of its formation is called the standard Gibbs energy of formation of this substance and is denoted ΔG 0. The relationship between ΔG and ΔG 0 is expressed by an equation called the Van't Hoff isotherm:

where R is the universal gas constant, T is temperature, K p is the equilibrium constant. For reaction

aA + bB = cC + dD

the equation can be written as:

or in the form:

For reactions occurring under isochoric-isothermal conditions, there is another criterion for the spontaneity of the process. Maximum work Wmax, which the system can accomplish during an equilibrium process under isochoric-isothermal conditions, is equal to the change Helmholtz energy systems ΔF (isochoric-isothermal potential):

ΔF = -W max.

The change in Helmholtz energy of the reaction is equal to

The change in Helmholtz energy characterizes the direction and limit of the spontaneous flow of a chemical reaction under isochoric-isothermal conditions, which is possible subject to the following inequality

ΔF< 0.

The relationship between thermodynamic functions is shown in Fig. 2.13.

ΔH ΔU pΔV TΔS ΔF TΔS ΔG

Rice. 2.13 Relationship between thermodynamic functions

Thermodynamic potentials are of great importance in determining the so-called chemical affinity . It has been experimentally established that some chemicals react with each other easily and quickly, others with difficulty, and others do not react at all. This gave rise to the introduction of the concept of chemical affinity, which can be defined as the ability of different substances to react with each other.

What is a measure of chemical affinity? The answer to this question turned out to be difficult. It was initially assumed that the rate of reaction between these substances could be taken as a measure of chemical affinity. But this assumption had to be abandoned, if only because the reaction rate depends not only on the chemical properties of the reagents and the parameters at which the reaction occurs, but also on the presence of catalysts - substances that do not participate in the reaction to any significant extent, but can very significantly influence its speed. The second assumption was that chemical affinity depends on the thermal effect of the reaction. But this assumption also did not stand up to scrutiny, since in different reactions the thermal effects have different signs.

Finally, it was found that a measure of chemical affinity is best determined by the change (decrease) in thermodynamic potential as a result of a reaction. Thus, thermodynamic potentials are of very great practical importance in chemistry. By calculating thermodynamic potentials, it is possible to determine measures of the chemical affinity of various substances, the possibility of carrying out a chemical reaction and its limits (equilibrium composition) depending on external conditions and, above all, on temperature.

Questions for self-control

1. What are the differences in the nature of the change in enthalpy of the system during an exo- and endothermic process.

2. How does the strength of chemical bonds in reaction products and starting materials affect the thermal effect of the reaction.

3. Formulate the concept of “enthalpy (heat) of formation of a substance”

4. a) Why are chemical and phase transformations accompanied by the release or absorption of energy?

5. Formulate the concept of heat capacity.

6. Formulate the I, II and III principles of thermodynamics.

Topic 3.

Related information.

Properties of entropy. Entropy - criterion of direction

1. Entropy is a function of the state of the system, i.e. its change D S depends only on the entropy of the initial and final states of the system.

2. Entropy characterizes the probability of the system’s implementation. The greater the entropy, the more ways the system can be implemented. For example, entropy increases when BMC molecules disintegrate into separate fragments, when a substance transitions from solid to liquid and gaseous state at a constant temperature, when the substance is heated (as the thermal movement of molecules increases and disorder increases). This relationship is expressed quantitatively by the Boltzmann formula

S = k ln W,

Where W – thermodynamic probability; k– Boltzmann constant, k= 1.38×10 -23 J/K.

Thermodynamic probability W is the number of microstates of the system with the help of which a given macrostate is realized. The macrostate of the system is characterized by state parameters ( p, V, T, chem. compound). But a thermodynamic system consists of a huge number of microparticles that have a certain energy, speed, and direction of movement, since they are in continuous chaotic motion. At equilibrium, the macrostate does not change, i.e. macro properties ( p, V, T, chem. composition) remain constant, but the microproperties (the position of the particle in the volume of the system, its energy, its speed) continuously change. The observed macrostate is realized by different microstates, the number of which is characterized by thermodynamic probability. In contrast to mathematical probability, which is equal to the ratio of the number of favorable events to the total number of possible events, and therefore is always less than unity, thermodynamic probability can be a very large value.

3.Entropy is a criterion for the direction of a spontaneous process in an isolated system.

In isolated systems there is no heat supply from the external environment ( Q= 0), therefore, according to the II law of thermodynamics (2), in an isolated system, entropy either remains constant in a state of equilibrium or increases during an irreversible (spontaneous) process. The growth of entropy continues until an equilibrium state is established, while the entropy value is maximum S max (picture).

Calculation of entropy changes during a phase transition,

heating (cooling), during a chemical reaction

For real (irreversible) processes, the II law of thermodynamics is written in the idea of ​​inequality, which makes it difficult to calculate the change in entropy D S during their course. But entropy is a function of the state of the system, and its change does not depend on the path of the process. Therefore, to calculate D S when various processes occur, we use the equation of the II law for reversible processes:

Entropy change during phase transformations

Phase transformation (phase transition)– a process associated with a change in the state of aggregation of a substance.

A characteristic feature of these processes is that they occur at a constant temperature - the phase transition temperature T f.p. .

Then, according to the II law of thermodynamics

Where Q f.p. . – thermal effect of phase transition.

At p = const heat is equal to the enthalpy change: