27. Structure and properties of iron; metastable and stable iron-carbon phase diagrams. Formation of the structure of carbon steels. Determination of the carbon content in steel by structure Iron-carbon alloys are the most common metal

Introduction.

Phases are called homogeneous different parts of physico-chemical systems. A substance is homogeneous when all the parameters of the state of the substance are the same in all its volumes, the dimensions of which are large compared to the interatomic states. Mixtures of different gases always form one phase if they are in the same concentration throughout the volume.

The same substance, depending on external conditions, can be in one of three states of aggregation - liquid, solid or gaseous. Depending on external conditions, it can be in one phase, or in several phases at once. In the nature around us, we especially often observe phase transitions of water. For example: evaporation, condensation. There are pressure and temperature conditions under which the substance is in equilibrium in different phases. For example, when liquefying a gas in a state of phase equilibrium, the volume can be anything, and the transition temperature is related to the saturation vapor pressure. The temperatures at which transitions from one phase to another occur are called transition temperatures. They depend on pressure, although to varying degrees: the melting point is weaker, the temperature of vaporization and sublimation is stronger. At normal and constant pressure, the transition occurs at a certain temperature, and here melting, boiling and sublimation (or sublimation.) take place. Sublimation is the transition of a substance from a solid to a gaseous state, which can be observed, for example, in the shells of cometary tails. When a comet is far from the sun, almost all of its mass is concentrated in its nucleus, which measures 10-12 kilometers. The nucleus surrounded by a small shell of gas is the so-called head of a comet. When approaching the Sun, the core and shells of the comet begin to heat up, the probability of sublimation increases, and desublimation decreases. The gases escaping from the nucleus of the comet entrain solid particles, the head of the comet increases in volume and becomes gas and dust in composition.

Phase transitions of the first and second kind.

Phase transitions are of several kinds. Changes in the aggregate states of a substance are called first-order phase transitions if:

1) The temperature is constant during the entire transition.

2) The volume of the system is changing.

3) The entropy of the system changes.

For such a phase transition to occur, it is necessary for a given mass of substance to sheathe a certain amount of heat corresponding to the latent heat of transformation. Indeed, during the transition of the condensed phase to a phase with a lower density, a certain amount of energy must be imparted in the form of heat, which will go to destroy the crystal lattice (during melting) or to remove liquid molecules from each other (during vaporization). During the transformation, latent heat will go to the transformation of cohesive forces, the intensity of thermal motion will not change, as a result, the temperature will remain constant. With such a transition, the degree of disorder, and hence the entropy, increases. If the process goes in the opposite direction, then latent heat is released. Phase transitions of the first kind include: the transformation of a solid into a liquid (melting) and the reverse process (crystallization), liquid into vapor (evaporation, boiling). One crystalline modification - to another (polymorphic transformations). Phase transitions of the second kind include: the transition of a normal conductor to a superconducting state, helium-1 to superfluid helium-2, a ferromagnet to a paramagnet. Metals such as iron, cobalt, nickel and gadolinium stand out for their ability to be highly magnetized and to maintain a state of magnetization for a long time. They are called ferromagnets. Most metals (alkali and alkaline earth metals and a significant part of transition metals) are weakly magnetized and do not retain this state outside a magnetic field - these are paramagnets. Phase transitions of the second, third, and so on kind are associated with the order of those derivatives of the thermodynamic potential? f, which experience finite measurements at the transition point. Such a classification of phase transformations is associated with the work of the theoretical physicist Paul Ernest (1880 -1933). So, in the case of a second-order phase transition, second-order derivatives experience jumps at the transition point: heat capacity at constant pressure ?p 2), thermal expansion coefficient b \u003d (1 / V 0) (? 2 f /? Tp), while the first derivatives remain continuous. This means that there is no release (absorption) of heat and no change in specific volume (φ - thermodynamic potential).

The state of phase equilibrium is characterized by a certain relationship between the phase transformation temperature and pressure. Numerically, this dependence for phase transitions is given by the Clausius-Clapeyron equation: p/T=q/TV. Research at low temperatures is a very important branch of physics. The fact is that in this way it is possible to get rid of interference associated with chaotic thermal motion and study phenomena in a “pure” form. This is especially important in the study of quantum regularities. Usually, due to chaotic thermal motion, a physical quantity is averaged over a large number of its different values, and quantum jumps are “smeared out”.

Low temperatures (cryogenic temperatures), in physics and cryogenic technology, the temperature range is below 120°K (0°C=273°K); the work of Carnot (he worked on a heat engine) and Clausius laid the foundation for research on the properties of gases and vapors, or technical thermodynamics. In 1850, Clausius noticed that saturated water vapor partially condenses during expansion and becomes superheated during compression. Renu made a special contribution to the development of this scientific discipline. The intrinsic volume of gas molecules at room temperature is approximately one thousandth of the volume occupied by the gas. In addition, molecules are attracted to each other at distances greater than those from which their repulsion begins.

A phase is a thermodynamically equilibrium state of a substance that differs in physical properties from other possible equilibrium states of the same substance. If, for example, there is water in a closed vessel, then this system is two-phase: liquid phase - water; gaseous phase - a mixture of air and water vapor. If pieces of ice are thrown into water, then this system will become three-phase, in which ice is a solid phase. Often the concept of "phase" is used in the sense of the state of aggregation, but it must be borne in mind that it is wider than the concept of "aggregate state". Within one state of aggregation, a substance can be in several phases that differ in their properties, composition and structure (ice, for example, occurs in five different modifications - phases). The transition of a substance from one phase to another - a phase transition - is always associated with qualitative changes in the properties of the substance. An example of a phase transition can be changes in the aggregate state of a substance or transitions associated with changes in the composition, structure and properties of a substance (for example, the transition of a crystalline substance from one modification to another).

There are two kinds of phase transitions. A phase transition of the first kind (for example, melting, crystallization, etc.) is accompanied by the absorption or release of heat, called the heat of the phase transition. Phase transitions of the first kind are characterized by the constancy of temperature, changes in entropy and volume. An explanation for this can be given as follows. For example, during melting, a certain amount of heat must be imparted to the body in order to cause the destruction of the crystal lattice. The heat supplied during melting is not used to heat the body, but to break interatomic bonds, so melting proceeds at a constant temperature. In such transitions - from a more ordered crystalline state to a less ordered liquid state - the degree of disorder increases, i.e., according to the second law of thermodynamics, this process is associated with an increase in the entropy of the system. If the transition occurs in the opposite direction (crystallization), then the system releases heat.

Phase transitions that are not associated with the absorption or release of heat and a change in volume are called phase transitions of the second order. These transitions are characterized by a constant volume and entropy, but an abrupt change in heat capacity. The general interpretation of phase transitions of the second kind was proposed by Academician L. D. Landau (1908-1968). According to this interpretation, second-order phase transitions are associated with a change in symmetry: above the transition point, the system, as a rule, has a higher symmetry than below the transition point. Examples of phase transitions of the second kind are: the transition of ferromagnetic substances (iron, nickel) at certain pressure and temperature to a paramagnetic state; the transition of metals and some alloys at a temperature close to 0 K into a superconducting state, characterized by an abrupt decrease in electrical resistance to zero; transformation of ordinary liquid helium (helium I) at T=2.9K into another liquid modification (helium II) with superfluidity properties.

Phases- these are various homogeneous parts of physico-chemical systems. A substance is homogeneous when all the parameters of the state of the substance are the same in all its elementary volumes, the dimensions of which are large compared to the interatomic states. Mixtures of different gases always form one phase if they are in the same concentration throughout the volume. The same substance, depending on external conditions, can be in one of three states of aggregation - liquid, solid or gaseous. Phases are stable states of a certain state of aggregation. The concept of a phase is broader than the concept of an aggregate state.

Depending on external conditions, the system can be in equilibrium either in one phase or in several phases at once. Their equilibrium existence is called phase balance.

Evaporation and condensation - frequently observed phase transitions of water in the natural environment. When water passes into steam, evaporation first occurs - the transition of the surface layer of the liquid into steam, while only the fastest molecules pass into steam: they must overcome the attraction of surrounding molecules, therefore their average kinetic energy and, accordingly, the temperature of the liquid decrease. Observed in everyday life and the reverse process - condensation. Both of these processes depend on external conditions. In some cases, a dynamic equilibrium is established between them, when the number of molecules leaving the liquid becomes equal to the number of molecules returning to it. Molecules in a liquid are bound by attractive forces that hold them within the liquid. If molecules with velocities that exceed the average are near the surface, they can leave it. Then the average speed of the remaining molecules will decrease and the temperature of the liquid will decrease. For evaporation at a constant temperature, a certain amount of heat must be imparted to the liquid: Q= rt, where r is the specific heat of vaporization, which decreases with increasing temperature. At room temperature, for one molecule of water, the heat of vaporization is 10 -20 J, while the average energy of thermal motion is 6.06 10 -21 J. This means that

molecules with an energy that is 10 times the energy of thermal motion. When passing through the liquid surface, the potential energy of a fast molecule increases, while the kinetic energy decreases. Therefore, the average kinetic energies of vapor and liquid molecules at thermal equilibrium are equal.

Saturated steam - it is a vapor in dynamic equilibrium, corresponding to a given temperature, with its liquid. Experience shows that it does not obey the Boyle-Mariotte law, since its pressure does not depend on volume. Saturated vapor pressure is the highest pressure that steam can have at a given temperature. The processes of evaporation and condensation of water cause complex interactions between the atmosphere and the hydrosphere, which are important for the formation of weather and climate. There is a continuous exchange of matter (water cycle) and energy between the atmosphere and the hydrosphere.

Studies have shown that about 7,000 km 3 of water evaporates per day from the surface of the World Ocean, which makes up 94% of the earth's hydrosphere, and about the same amount falls in the form of precipitation. Water vapor, carried away by the convection movement of air, rises up and enters the cold layers of the troposphere. As it rises, the vapor becomes more and more saturated, then condenses to form raindrops. In the process of steam condensation in the troposphere, about 1.6-10 22 J of heat is released per day, which is tens of thousands of times greater than the energy generated by mankind over the same time.

Boiling- the process of transition of a liquid into vapor as a result of the emergence of bubbles filled with vapor. Boiling occurs throughout the volume. The rupture of bubbles at the surface of a boiling liquid indicates that the vapor pressure in them exceeds the pressure above the surface of the liquid. At a temperature of 100 °C, the saturated vapor pressure is equal to the air pressure above the surface of the liquid (this is how this point on the scale was chosen). At an altitude of 5 km, the air pressure is half as much and water boils there at 82 ° C, and at the border of the troposphere (17 km) - at approximately 65 ° C. Therefore, the boiling point of a liquid corresponds to the temperature at which its saturated vapor pressure is equal to the external pressure. The weak gravitational field of the Moon (gravitational acceleration near its surface is only 1.7 m/s 2) is not able to hold the atmosphere, and in the absence of atmospheric pressure, the liquid instantly boils away, so the lunar "seas" are waterless and are formed by solidified lava. For the same reason, the Martian "channels" are also waterless.

A substance can be in equilibrium and in different phases. So, when liquefying a gas in a state of phase equilibrium, the volume can be anything, and the transition temperature is related to the saturation vapor pressure. The phase equilibrium curve can be obtained by projecting onto a plane (p, t) areas of transition to the liquid state. Analytically, the equilibrium curve of two phases is determined from the solution of the Clausius-Clapeyron differential equation. Similarly, it is possible to obtain melting and sublimation curves, which are connected at one point of the plane (R, D), at the triple point (see Fig. 7.1), where in certain proportions they are in equal

all three phases. The triple point of water corresponds to a pressure of 569.24 Pa and a temperature of -0.0075 °C; carbon dioxide - 5.18 10 5 Pa and 56.6 ° C, respectively. Therefore, at atmospheric pressure R, equal to 101.3 kPa, carbon dioxide can be in a solid or gaseous state. At the critical temperature, the physical properties of liquid and vapor become the same. At temperatures above the critical point, the substance can only be in the gaseous state. For water - T= 374.2 °С, R= 22.12 MPa; for chlorine - 144 ° C and 7.71 MPa, respectively.

Transition temperatures are the temperatures at which transitions from one phase to another occur. They depend on pressure, although to varying degrees: the melting point is weaker, the temperatures of vaporization and sublimation are stronger. At normal and constant pressures, the transition occurs at a certain temperature, and here melting, boiling and sublimation (or sublimation) points take place.

The transition of matter from a solid state directly to a gaseous state can be observed, for example, in the shells of cometary tails. When a comet is far from the Sun, almost all of its mass is concentrated in its nucleus, which has a size of 10-12 km. The nucleus is surrounded by a small shell of gas - this is the comet's head. When approaching the Sun, the core and shell of the comet begin to heat up, the probability of sublimation increases, and desublimation (the reverse process) decreases. The gases escaping from the comet's nucleus carry away solid particles, the comet's head increases in volume and becomes gas and dust in composition. The pressure of the cometary nucleus is very low, so the liquid phase does not occur. Along with the head, the comet's tail also grows, which stretches away from the Sun. In some comets it reaches hundreds of millions of kilometers at perihelion, but the densities in the cometary matter are negligible. With each approach to the Sun, comets lose most of their mass, more and more volatile substances sublimate in the nucleus, and gradually it crumbles into meteor bodies that form meteor showers. Over the 5 billion years of the existence of the solar system, many comets ended their existence this way.

In the spring of 1986, automatic Soviet stations "Vega-1" and "Vega-2" were sent into space to study Halley's comet, which passed at a distance of 9000 and 8200 km from it, respectively, and the NASA station "Giotto" - at a distance of only 600 km from the comet's nucleus. The nucleus was 14 x 7.5 km in size, dark in color and about 400 K in temperature. When the space stations passed through the comet's head, about 40,000 kg of icy matter sublimated in 1 s.

In late autumn, when a sharp cold snap sets in after wet weather, one can observe on the branches of trees and on wires

frost is desublimated ice crystals. A similar phenomenon is used when storing ice cream, when carbon dioxide is cooled, as the molecules passing into steam carry away energy. On Mars, the phenomena of sublimation and desublimation of carbon dioxide in the polar caps play the same role as evaporation - condensation in the atmosphere and hydrosphere of the Earth.

The heat capacity tends to zero at ultralow temperatures, as Nernst established. From this, Planck showed that near absolute zero, all processes proceed without a change in entropy. Einstein's theory of the heat capacity of solids at low temperatures made it possible to formulate Nernst's result as the third law of thermodynamics. The unusual properties of substances observed at low temperatures - superfluidity and superconductivity - have been explained in modern theory as macroscopic quantum effects.

Phase transitions are of several kinds. During a phase transition, the temperature does not change, but the volume of the system does.

Phase transitions of the first kind changes in the aggregate states of a substance are called if: the temperature is constant during the entire transition; the volume of the system changes; the entropy of the system changes. For such a phase transition to occur, it is necessary to impart a certain amount of heat to a given mass of substance, corresponding to the latent heat of transformation.

Indeed, during the transition from a more condensed phase to a phase with a lower density, a certain amount of energy must be imparted in the form of heat, which will go to destroy the crystal lattice (during melting) or to remove liquid molecules from each other (during vaporization). During the transformation, latent heat is expended to overcome cohesive forces, the intensity of thermal motion does not change, as a result, the temperature remains constant. With such a transition, the degree of disorder, and hence the entropy, increases. If the process goes in the opposite direction, then latent heat is released.

Phase transitions of the second kind associated with a change in the symmetry of the system: above the transition point, the system, as a rule, has a higher symmetry, as L.D. Landau showed in 1937. For example, in a magnet, the spin moments above the transition point are randomly oriented, and the simultaneous rotation of all spins around the same axis by the same angle does not change the properties of the system. Below the transition points, the spins have some preferential orientation, and their simultaneous rotation changes the direction of the magnetic moment of the system. Landau introduced the ordering factor and expanded the thermodynamic potential at the transition point in powers of this coefficient, on the basis of which he built a classification of all possible types of transitions.

Dov, as well as the theory of the phenomena of superfluidity and superconductivity. On this basis, Landau and Lifshitz considered many important problems - the transition of a ferroelectric to a paraelectric, a ferromagnet to a paramagnet, sound absorption at the transition point, the transition of metals and alloys to the superconducting state, etc.

The calculation of the thermodynamic properties of a system based on statistical mechanics involves the choice of a specific model of the system, and the more complex the system, the simpler the model should be. E. Ising proposed a model of a ferromagnet (1925) and solved the problem of a one-dimensional chain, taking into account the interaction with the nearest neighbors for any fields and temperatures. In the mathematical description of such systems of particles with intense interaction, a simplified model is chosen, when only pair-type interaction occurs (such a two-dimensional model is called the Ising lattice). But phase transitions were not always calculated, probably due to some unaccounted phenomena common to systems of many particles, and the nature of the particles themselves (liquid particles or magnets) does not matter. L. Onsager gave an exact solution for the two-dimensional Ising model (1944). He placed dipoles at the lattice nodes, which can orient themselves in only two ways, and each such dipole can only interact with its neighbor. It turned out that at the transition point, the heat capacity goes to infinity according to the logarithmic law symmetrically on both sides of the transition point. Later it turned out that this conclusion is very important for all second-order phase transitions. Onsager's work showed that the method of statistical mechanics makes it possible to obtain new results for phase transformations.

Phase transitions of the second, third, etc. genera are related to the order of those derivatives of the thermodynamic potential Ф, which experience finite changes at the transition point. Such a classification of phase transformations is associated with the work of the theoretical physicist P. Ehrenfest. In the case of a second-order phase transition, the second-order derivatives experience jumps at the transition point: heat capacity at constant pressure C p =, compressibility , coefficient

thermal expansion coefficient, while per-

all derivatives remain continuous. This means that there is no release (absorption) of heat and no change in specific volume.

Quantum field theory began to be used for calculations of particle systems only in the 70s. 20th century The system was considered as a lattice with a variable step, which made it possible to change the accuracy of calculations and approach the description of a real system and use a computer. The American theoretical physicist C. Wilson, having applied a new method of calculations, received a qualitative leap in understanding second-order phase transitions associated with rearrangement of the system's symmetry. In fact, he connected quantum mechanics with statistical, and his work received fundamental

mental meaning. They are applicable in combustion processes, and in electronics, and in the description of cosmic phenomena and nuclear interactions. Wilson investigated a wide class of critical phenomena and created a general theory of second-order phase transitions.

An important branch of thermodynamics is the study of transformations between different phases of a substance, since these processes occur in practice and are of fundamental importance for predicting the behavior of a system under certain conditions. These transformations are called phase transitions, to which the article is dedicated.

The concept of a phase and a system component

Before proceeding to the consideration of phase transitions in physics, it is necessary to define the concept of the phase itself. As is known from the course of general physics, there are three states of matter: gaseous, solid and liquid. In a special section of science - in thermodynamics - the laws are formulated for the phases of matter, and not for their states of aggregation. A phase is understood as a certain volume of matter that has a homogeneous structure, is characterized by specific physical and chemical properties and is separated from the rest of the matter by boundaries, which are called interphase.

Thus, the concept of "phase" carries much more practically significant information about the properties of matter than its state of aggregation. For example, the solid state of a metal such as iron may be in the following phases: low temperature magnetic body centered cubic (BCC), low temperature nonmagnetic bcc, face centered cubic (fcc), and high temperature nonmagnetic bcc.

In addition to the concept of "phase", the laws of thermodynamics also use the term "components", which means the number of chemical elements that make up a particular system. This means that the phase can be both monocomponent (1 chemical element) and multicomponent (several chemical elements).

Gibbs' theorem and equilibrium between phases of a system

To understand phase transitions, it is necessary to know the equilibrium conditions between them. These conditions can be mathematically obtained by solving the system of Gibbs equations for each of them, assuming that the equilibrium state is reached when the total Gibbs energy of the system isolated from external influence ceases to change.

As a result of solving this system of equations, conditions are obtained for the existence of equilibrium between several phases: an isolated system will cease to evolve only when the pressures, chemical potentials of each component and temperatures in all phases are equal to each other.

Gibbs phase rule for equilibrium

A system consisting of several phases and components can be in equilibrium not only under certain conditions, for example, at a specific temperature and pressure. Some of the variables in the Gibbs theorem for equilibrium can be changed while maintaining both the number of phases and the number of components that are in this equilibrium. The number of variables that can be changed without disturbing the equilibrium in the system is called the number of freedoms of this system.

The number of freedoms l of a system consisting of f phases and k components is uniquely determined from the Gibbs phase rule. This rule is mathematically written as follows: l + f = k + 2. How to work with this rule? Very simple. For example, it is known that the system consists of f=3 equilibrium phases. What is the minimum number of components such a system can contain? You can answer the question by reasoning as follows: in the case of equilibrium, the most stringent conditions exist when it is realized only at certain indicators, that is, a change in any thermodynamic parameter will lead to imbalance. This means that the number of freedoms l=0. Substituting the known values ​​of l and f, we obtain k=1, that is, a system in which three phases are in equilibrium can consist of one component. A striking example is the triple point of water, when ice, liquid water and steam exist in equilibrium at specific temperatures and pressures.

Classification of phase transformations

If you begin to change some in a system that is in equilibrium, then you can observe how one phase will disappear, and another will appear. A simple example of this process is the melting of ice when it is heated.

Given that the Gibbs equation depends on only two variables (pressure and temperature), and a phase transition involves a change in these variables, then mathematically the transition between phases can be described by differentiating the Gibbs energy with respect to its variables. It was this approach that was used by the Austrian physicist Paul Ehrenfest in 1933, when he compiled a classification of all known thermodynamic processes that occur with a change in phase equilibrium.

It follows from the basics of thermodynamics that the first derivative of the Gibbs energy with respect to temperature is equal to the change in the entropy of the system. The derivative of the Gibbs energy with respect to pressure is equal to the change in volume. If, when the phases in the system change, the entropy or volume suffers a break, that is, they change sharply, then they speak of a first-order phase transition.

Further, the second derivatives of the Gibbs energy with respect to temperature and pressure are the heat capacity and the coefficient of volumetric expansion, respectively. If the transformation between phases is accompanied by a discontinuity in the values ​​of the indicated physical quantities, then one speaks of a second-order phase transition.

Examples of transformations between phases

There are a huge number of different transitions in nature. Within the framework of this classification, striking examples of transitions of the first kind are the processes of melting metals or the condensation of water vapor from air, when there is a volume jump in the system.

If we talk about transitions of the second kind, then striking examples are the transformation of iron from a magnetic to a paramagnetic state at a temperature of 768 ºC or the transformation of a metallic conductor into a superconducting state at temperatures close to absolute zero.

Equations that describe transitions of the first kind

In practice, it is often necessary to know how the temperature, pressure, and absorbed (released) energy change in a system when phase transformations occur in it. Two important equations are used for this purpose. They are obtained based on knowledge of the basics of thermodynamics:

1. Clapeyron's formula, which establishes the relationship between pressure and temperature during transformations between different phases.
2. The Clausius formula, which relates the absorbed (released) energy and the temperature of the system during the transformation.

The use of both equations is not only in obtaining quantitative dependences of physical quantities, but also in determining the sign of the slope of equilibrium curves in phase diagrams.

Equation for describing transitions of the second kind

Phase transitions of the 1st and 2nd kind are described by different equations, since the use of and Clausius for transitions of the second kind leads to mathematical uncertainty.

To describe the latter, the Ehrenfest equations are used, which establish a relationship between changes in pressure and temperature through knowledge of the change in heat capacity and the coefficient of volumetric expansion during the transformation process. The Ehrenfest equations are used to describe conductor-superconductor transitions in the absence of a magnetic field.

Importance of phase diagrams

Phase diagrams are a graphic representation of areas in which the corresponding phases exist in equilibrium. These areas are separated by equilibrium lines between the phases. P-T (pressure-temperature), T-V (temperature-volume) and P-V (pressure-volume) axes are often used.

The importance of phase diagrams lies in the fact that they allow you to predict what phase the system will be in when the external conditions change accordingly. This information is used in the heat treatment of various materials in order to obtain a structure with desired properties.