A liquid, such as water, can be in solid, liquid and gaseous states, which are called phase states of matter. In liquids, the distances between molecules are approximately two orders of magnitude smaller than in gases. In a solid, the molecules are located even closer to each other. The temperature at which it changes phase state of matter(liquid - solid, liquid - gaseous), called phase transition temperature.

Heat of phase transition or latent heat is the amount of heat of melting or evaporation of a substance. Figure 6.9 shows the dependence of water temperature on the amount of heat received in calories. It can be seen that at temperatures 0 0 C and 100 0 C, a change in the phase state of water occurs, but the temperature of the water does not change. The absorbed heat is spent on changing the phase state of the substance. Physically, this means that when a solid body, for example ice, is heated at 0 0 C, the amplitude of vibrations of the molecules relative to each other increases. This leads to an increase in their potential energy, and, consequently, to a weakening or breaking of intermolecular bonds. Molecules or their clusters are able to move relative to each other. Ice turns into liquid at a constant temperature. After changing its state of aggregation from solid to liquid, the absorption of heat leads to an increase in temperature according to a linear law. This happens up to 100 0 C. Then the energy of the vibrating molecules increases so much that the molecules are able to overcome the attraction of other molecules. They violently break away not only from the surface of the water, but also form bubbles of steam throughout the entire volume of the liquid. They rise to the surface under the action of buoyant force and are thrown out. In this phase change, water turns into steam. Further, the absorption of heat again leads to an increase in the temperature of the steam according to a linear law.

The heat released or absorbed during a phase transition depends on the mass of the substance.

When a substance of mass m passes from liquid to gaseous state or, conversely, from gaseous to liquid, heat Q is absorbed or released:

Specific heat of vaporization Q required to convert 1 kg of liquid into steam at boiling point:

When a substance passes from solid to liquid and back, an amount of heat is absorbed or transferred:

Specific heat of fusion q is called the amount of heat Q required to convert 1 kg of solid (e.g. ice) into liquid at melting point:

The specific heat of fusion and vaporization is measured in J/kg. With increasing temperature, the specific heat of vaporization decreases, and at a critical temperature it becomes equal to zero.

For water, the specific heat of fusion and vaporization are respectively:

, .

Here, an off-system unit of measurement of the amount of energy is used - the calorie, equal to the amount of heat required to heat 1 gram of water by 1 °C at normal atmospheric pressure of 101.325 kPa.

As can be seen in Fig. 6.17, heating ice from -20 0 C to 0 0 C requires eight times less energy than turning it from ice into water, and 54 times less than turning water into steam.

Fig.6.17. Dependence of temperature on heat supplied to the system

for 1 kg of ice.

The temperature at which the difference between vapor and liquid is lost is called critical. In Fig. Figure 6.18 illustrates the concept of critical temperature on the dependence of the density of water and steam on temperature. When water is heated in a closed test tube, as can be seen in Fig. 6.18, the density of water decreases with increasing temperature due to the volumetric expansion of water, and the density of steam increases. At a certain temperature, which is called critical, the density of steam becomes equal to the density of water.

Each substance has its own critical temperature. For water, nitrogen and helium, the critical temperatures are respectively:

, , .

Fig.6.18. Critical point on the dependence graph

density of steam and water on temperature.

Fig.6.19. Dependence of pressure on volume p=p(V) for steam. In the area indicated by the dotted line, the gaseous and liquid states of the substance exist simultaneously.

Figure 6.19 shows the dependence of steam pressure on its volume P=P(V). The equation of state of steam at low pressure and far from the temperature of its phase transition (above point b 0 in Fig. 6.19) is close to the equation of state of an ideal gas (that is, in this case the gas can be considered ideal and its behavior is well described by the Boyle-Moriott law). As the temperature decreases, the dependence P=P(V) begins to differ from its form for an ideal gas. Location on a – b steam condensation occurs and the steam pressure remains almost unchanged, and the dependence in Fig. 6.19 is a slowly decreasing linear function. Below the point A, all the vapor becomes a liquid, and then compression of the liquid occurs. In this case, as can be seen in Fig. 6.11, the pressure with a very slight decrease in volume, since the liquid is practically incompressible, increases sharply.

Since the temperature of the phase transition depends on the gas pressure, phase transitions can be represented using the dependence of pressure on temperature P=P(T) in Fig. 6.20. A change in the phase state of a substance occurs at the boundary of vapor - liquid, solid - liquid, solid - vapor. On different sides of these boundary lines, the gas is in a different state of aggregation - solid, liquid or gaseous.

Fig.6.20. Phase diagram for water.

The point of intersection of three lines in Fig. 6.12 is called triple point. For example, water at a temperature of 0 0 C and a pressure of atm. has a triple point, and carbon dioxide has a triple point at a temperature and pressure of P = 5.1 atm. Figure 6.20 shows that it is possible for a substance to transition from a gaseous to a solid state and vice versa, bypassing the liquid stage.

The transition from the solid state of a substance to the gaseous state is called sublimation.

Example: cooling with dry ice, for example, packs of ice cream on trays. In this case, as we have seen many times, dry ice turns into steam.

How to turn gas into liquid? The boiling point chart answers this question. You can turn a gas into a liquid by either decreasing the temperature or increasing the pressure.

In the 19th century, increasing pressure seemed an easier task than lowering temperature. At the beginning of this century, the great English physicist Michael Farada managed to compress gases to vapor pressure values ​​and in this way turn many gases (chlorine, carbon dioxide, etc.) into liquid.

However, some gases - hydrogen, nitrogen, oxygen - could not be liquefied. No matter how much pressure was increased, they did not turn into liquid. One might think that oxygen and other gases cannot be liquid. They were classified as true, or permanent, gases.

In fact, the failures were caused by a lack of understanding of one important circumstance.

Let's consider liquid and vapor in equilibrium and think about what happens to them as the boiling point increases and, of course, the corresponding increase in pressure. In other words, imagine that a point on the boiling graph moves upward along the curve. It is clear that as the temperature increases, a liquid expands and its density decreases. As for steam, does the boiling point increase? Of course, it contributes to its expansion, but, as we have already said, the saturated vapor pressure increases much faster than the boiling point. Therefore, the vapor density does not fall, but, on the contrary, quickly increases with increasing boiling temperature.

Since the density of the liquid decreases and the density of the vapor increases, then, moving “up” along the boiling curve, we will inevitably reach a point at which the densities of the liquid and vapor are equal (Fig. 4.3).

At this remarkable point, called the critical point, the boiling curve ends. Since all the differences between gas and liquid are associated with the difference in density, at the critical point the properties of the liquid and gas become the same. Each substance has its own critical temperature and its own critical pressure. Thus, for water, the critical point corresponds to a temperature of 374 ° C and a pressure of 218.5 atm.

If you compress a gas whose temperature is below the critical temperature, then the process of its compression will be represented by an arrow crossing the boiling curve (Fig. 4.4). This means that at the moment of reaching a pressure equal to the vapor pressure (the point where the arrow intersects the boiling curve), the gas will begin to condense into a liquid. If our vessel were transparent, then at this moment we would see the beginning of the formation of a layer of liquid at the bottom of the vessel. At constant pressure, the layer of liquid will grow until finally all the gas turns into liquid. Further compression will require an increase in pressure.

The situation is completely different when compressing a gas whose temperature is above the critical temperature. The compression process can again be depicted as an arrow going from bottom to top. But now this arrow does not cross the boiling curve. This means that when compressed, the steam will not condense, but will only be continuously compacted.

At temperatures above the critical temperature, the existence of liquid and gas separated by an interface is impossible: When compressed to any density, there will be a homogeneous substance under the piston, and it is difficult to say when it can be called a gas and when a liquid.

The presence of a critical point shows that there is no fundamental difference between the liquid and gaseous states. At first glance, it might seem that there is no such fundamental difference only when we are talking about temperatures above the critical one. This, however, is not the case. The existence of a critical point indicates the possibility of turning a liquid - a real liquid that can be poured into a glass - into a gaseous state without any semblance of boiling.

This transformation path is shown in Fig. 4.4. A cross marks a known liquid. If you lower the pressure a little (down arrow), it will boil, and it will also boil if you raise the temperature a little (arrow to the right). But we will do something completely different. We will compress the liquid very strongly, to a pressure above critical. The point representing the state of the liquid will go vertically upward. Then we heat the liquid - this process is depicted by a horizontal line. Now, after we find ourselves to the right of the Critical Temperature, we lower the pressure to the original one. If you now reduce the temperature, you can get real steam, which could be obtained from this liquid in a simpler and shorter way.

Thus, it is always possible, by changing pressure and temperature bypassing the critical point, to obtain steam by continuously transferring it from liquid or liquid from steam. This continuous transition does not require boiling or condensation.

Early attempts to liquefy gases such as oxygen, nitrogen, and hydrogen were unsuccessful because the existence of a critical temperature was not known. These gases have very low critical temperatures: nitrogen -147°C, oxygen -119°C, hydrogen -240°C, or 33 K. The record holder is helium, its critical temperature is 4.3 K. Convert these gases into there is only one way to liquid - you need to reduce their temperature below the specified"

The phase equilibrium curve (in the P, T plane) may end at some point (Fig. 16); such a point is called critical, and the temperature and pressure corresponding to it are called critical temperature and critical pressure. At higher temperatures and at higher pressures, there are no different phases, and the body is always homogeneous.

We can say that at the critical point the difference between both phases disappears. The concept of a critical point was first introduced by D.I. Mendeleev (1860).

In coordinates T, V, the equilibrium diagram in the presence of a critical point looks as shown in Fig. 17. As the temperature approaches its critical value, the specific volumes of the phases in equilibrium with each other come closer and coincide at the critical point (K in Fig. 17). The diagram in coordinates P, V has a similar appearance.

In the presence of a critical point, a continuous transition can be made between any two states of a substance, in which at no moment does separation into two phases occur - for this it is necessary to change the state along some curve that envelops the critical point and does not intersect the equilibrium curve anywhere. In this sense, in the presence of a critical point, the very concept of different phases becomes conditional, and it is impossible in all cases to indicate which states are one phase and which are another. Strictly speaking, we can talk about two phases only when they both exist simultaneously, touching each other, that is, at points lying on the equilibrium curve.

It is clear that the critical point can exist only for such phases, the difference between which is only of a purely quantitative nature. These are liquid and gas, differing from each other only in the greater or lesser role of the interaction between molecules.

The same phases as liquid and solid (crystal) or various crystalline modifications of a substance are qualitatively different from each other, since they differ in their internal symmetry. It is clear that about any property (element) of symmetry one can only say either that it exists or that it does not exist; it can appear or disappear only immediately, abruptly, and not gradually. In each state the body will have either one or the other symmetry, and therefore it is always possible to indicate which of the two phases it belongs to. The critical point, therefore, cannot exist for such phases, and the equilibrium curve must either go to infinity or end by intersecting the equilibrium curves of other phases.

The usual point of phase transition does not represent, in mathematical terms, singularities for the thermodynamic quantities of a substance. Indeed, each of the phases can exist (at least as metastable) on the other side of the transition point; thermodynamic inequalities are not violated at this point. At the transition point, the chemical potentials of both phases are equal to each other: ; for each of the functions this point is not remarkable in any way.

Let us depict in the plane P, V any isotherm of liquid and gas, i.e., the curve of dependence of P on V during isothermal expansion of a homogeneous body in Fig. 18). According to the thermodynamic inequality, there is a decreasing function V. Such a slope of the isotherms should remain for some extent beyond the points of their intersection with the equilibrium curve of liquid and gas (points b and sections of the isotherms correspond to metastable superheated liquid and supercooled steam, in which thermodynamic inequalities are still observed ( Of course, the horizontal segment on which separation into two phases does not correspond to a completely equilibrium isothermal change of state between points b).

If we take into account that the points have the same ordinate P, then it is clear that both parts of the isotherm cannot transform into each other in a continuous manner, and there must be a gap between them. The isotherms end at points (c and d), at which the thermodynamic inequality is violated, i.e.

By constructing the geometric locus of the end points of the liquid and gas isotherms, we obtain the AKB curve, on which thermodynamic inequalities are violated (for a homogeneous body); it limits the region in which the body under no circumstances can exist as homogeneous. The areas between this curve and the phase equilibrium curve correspond to superheated liquid and supercooled steam. Obviously, at the critical point both curves must touch each other. Of the points lying on the AKB curve itself, only the critical point K corresponds to the actually existing states of a homogeneous body - the only one at which this curve comes into contact with the region of stable homogeneous states.

In contrast to the usual points of phase equilibrium, the critical point is, in mathematical terms, a special point for the thermodynamic functions of matter (the same applies to the entire AKB curve, which limits the region of existence of homogeneous states of the body). The nature of this feature and the behavior of matter near the critical point will be discussed in § 153.

Supercritical state- the fourth form of the state of aggregation into which many organic and inorganic substances can transform.

The supercritical state of matter was first discovered by Cagniard de la Tour in 1822. Real interest in the new phenomenon arose in 1869 after the experiments of T. Andrews. Conducting experiments in thick-walled glass tubes, the scientist investigated the properties CO2, which liquefies easily when pressure increases. As a result, he found that at 31 ° C and 7.2 MPa, the meniscus, the boundary separating the liquid and the vapor in equilibrium with it, disappears, while the system becomes homogeneous (homogeneous) and the entire volume takes on the appearance of a milky-white opalescent liquid. With a further increase in temperature, it quickly becomes transparent and mobile, consisting of constantly flowing jets, reminiscent of flows of warm air over a heated surface. Further increases in temperature and pressure did not lead to visible changes.

He called the point at which such a transition occurs critical, and the state of the substance located above this point - supercritical. Despite the fact that outwardly this state resembles a liquid, a special term is now used to apply to it - supercritical fluid (from the English word fluid, that is, “capable of flow”). In modern literature, the abbreviated designation for supercritical fluids is SCF.

The location of the lines delimiting the regions of gaseous, liquid and solid states, as well as the position of the triple point where all three regions converge, are individual for each substance. The supercritical region begins at the critical point (indicated by an asterisk), which is certainly characterized by two parameters - temperature ( T cr.) and pressure ( R cr.). A decrease in either temperature or pressure below critical values ​​removes the substance from the supercritical state.

The fact of the existence of a critical point made it possible to understand why some gases, for example, hydrogen, nitrogen and oxygen, for a long time could not be obtained in liquid form with increasing pressure, which is why they were called permanent gases (from the Latin permanentis- "constant"). The diagram above shows that the region of existence of the liquid phase is located to the left of the critical temperature line. Thus, to liquefy any gas, it must first be cooled to below a critical temperature. U CO 2 the critical temperature is above room temperature, so it can be liquefied under specified conditions by increasing the pressure. Nitrogen has a much lower critical temperature: -146.95 ° C, therefore, if you compress nitrogen under normal conditions, you can ultimately reach the supercritical region, but liquid nitrogen cannot be formed. It is necessary to first cool the nitrogen below a critical temperature and then, by increasing the pressure, reach the region where liquid existence is possible. The situation is similar for hydrogen and oxygen, so before liquefaction they are cooled to a temperature below critical, and only then the pressure is increased. The supercritical state is possible for most substances; it is only necessary that the substance does not decompose at a critical temperature. In comparison with these substances, the critical point of water is reached with great difficulty: t cr= 374.2° C and R cr = 21,4 MPa.

The critical point is recognized as an important physical parameter of a substance, the same as melting or boiling points. The density of SCF is extremely low; for example, water in the SCF state has a density three times lower than under normal conditions. All SCFs have extremely low viscosity.

Supercritical fluids are a cross between a liquid and a gas. They can be compressed like gases (ordinary liquids are practically incompressible) and, at the same time, are capable of dissolving many substances in solid and liquid states, which is unusual for gases. Supercritical ethanol (at temperatures above 234 ° C) very easily dissolves some inorganic salts ( CoCl2, KBr, KI). Carbon dioxide, nitrous oxide, ethylene and some other gases in the SCF state acquire the ability to dissolve many organic substances - stearic acid, paraffin, naphthalene. Properties of supercritical CO 2 As a solvent, it can be regulated - with increasing pressure, its dissolving ability increases sharply.

Supercritical fluids became widely used only in the 1980s, when the general level of industrial development made SCF plants widely available. From that moment on, the intensive development of supercritical technologies began. SCF are not only good solvents, but also substances with a high diffusion coefficient, i.e. they easily penetrate into the deep layers of various solids and materials. The most widely used is supercritical CO 2, which turned out to be a solvent for a wide range of organic compounds. Carbon dioxide has become a leader in the world of supercritical technologies because... has a whole range of advantages. It is quite easy to transfer it to a supercritical state ( t cr– 31° C, R cr – 73,8 atm.), in addition, it is non-toxic, non-flammable, non-explosive, moreover, it is cheap and available. From the point of view of any technologist, it is an ideal component of any process. What makes it particularly attractive is that it is an integral part of atmospheric air and, therefore, does not pollute the environment. Supercritical CO 2 can be considered an environmentally friendly, absolutely clean solvent. Let's give just a few examples of its use.

Caffeine, a drug used to improve the functioning of the cardiovascular system, is obtained from coffee beans even without first grinding them. Complete extraction is achieved due to the high penetrating ability of SCF. The grains are placed in an autoclave - a container that can withstand high pressure, then gaseous CO 2, then create the required pressure (>73 atm.), as a result CO 2 goes into a supercritical state. All contents are mixed, after which the fluid along with dissolved caffeine is poured into an open container. Carbon dioxide, once exposed to atmospheric pressure, turns into a gas and escapes into the atmosphere, while the extracted caffeine remains in an open container in its pure form.

The use of SCF has proven to be very successful in cleaning contaminants from electronic circuits during their manufacturing process, since no traces of cleaning solvent remain on them.

Due to the rapid pace of production of the active part of light oil reserves, interest in methods for increasing oil recovery has sharply increased. If in the 70–80s of the 20th century the number of projects aimed at solving the problem of increasing oil recovery by injecting miscible hydrocarbon solvents, “inert” gases and carbon dioxide was comparable, then at the end of the 20th and beginning of the 21st centuries only the injection method CO 2 had a steady growth trend. Efficiency of application CO 2 for increasing oil recovery has been proven not only by experimental and theoretical work, but also by the results of numerous industrial tests.

Do not forget that the technology of enhanced oil recovery using CO 2 allows us to simultaneously solve the problem of conserving the huge amount of carbon dioxide released by industry.

Features of the process of exposure to injected CO2 on an oil and gas deposit depend on its state of aggregation.

The excess of pressure and temperature above the critical values ​​for carbon dioxide (and this is the most likely situation in reservoir conditions) predetermines its supercritical state. In this case CO2, which has exceptional dissolving ability in relation to hydrocarbon liquids when directly dissolved in reservoir oil, reduces its viscosity and dramatically improves filtration properties. This circumstance gives every reason to classify SCF – enhanced oil recovery technologies – as one of the most promising.

CHAPTER IV.
THERMODYNAMICS OF SOLUTIONS (SOLUTIONS)

If a certain amount of liquid is placed in a closed vessel, then part of the liquid will evaporate and saturated steam will exist above the liquid. The pressure, and therefore the density of this vapor, depends on the temperature. The density of vapor is usually much less than the density of liquid at the same temperature. If you increase the temperature, the density of the liquid will decrease (§ 198), while the pressure and density of the saturated vapor will increase. In table Figure 22 shows the density values ​​of water and saturated water vapor for different temperatures (and therefore for the corresponding pressures). In Fig. 497 the same data is presented in graph form. The top part of the graph shows the change in the density of a liquid depending on its temperature. As the temperature increases, the density of the liquid decreases. The lower part of the graph shows the dependence of saturated vapor density on temperature. The vapor density increases. At the temperature corresponding to point , the densities of the liquid and saturated vapor coincide.

Rice. 497. Dependence of the density of water and its saturated vapor on temperature

Table 22. Properties of water and its saturated steam at different temperatures

The table shows that the higher the temperature, the smaller the difference between the density of the liquid and the density of its saturated vapor. At a certain temperature (at water) these densities coincide. The temperature at which the densities of the liquid and its saturated vapor coincide is called the critical temperature of the substance. In Fig. 497 corresponds to the dot. The pressure corresponding to point is called critical pressure. The critical temperatures of different substances vary greatly. Some of them are given in table. 23.

Table 23. Critical temperature and critical pressure of some substances

What does the existence of a critical temperature indicate? What happens at even higher temperatures?

Experience shows that at temperatures higher than critical, a substance can only be in a gaseous state. If we reduce the volume occupied by steam at a temperature above the critical temperature, then the pressure of the steam increases, but it does not become saturated and continues to remain homogeneous: no matter how high the pressure, we will not find two states separated by a sharp boundary, as is always observed at lower temperatures due to steam condensation. So, if the temperature of a substance is above the critical temperature, then equilibrium of the substance in the form of a liquid and the vapor in contact with it is impossible at any pressure.

The critical state of a substance can be observed using the device shown in Fig. 498. It consists of an iron box with windows, which can be heated higher (“air bath”), and a glass ampoule with ether located inside the bath. When the bath is heated, the meniscus in the ampoule rises, becomes flatter and finally disappears, which indicates a transition through a critical state. As the bath cools, the ampoule suddenly becomes cloudy due to the formation of many tiny droplets of ether, after which the ether collects at the bottom of the ampoule.

Rice. 498. Device for observing the critical state of the ether

As can be seen from table. 22, as the critical point is approached, the specific heat of vaporization becomes less and less. This is explained by the fact that as the temperature increases, the difference in the internal energies of a substance in the liquid and vapor states decreases. In fact, the adhesive forces of molecules depend on the distances between molecules. If the densities of liquid and vapor differ little, then the average distances between molecules differ little. Consequently, the values ​​of the potential energy of interaction between molecules will differ little. The second term of the heat of vaporization - work against external pressure - also decreases as the critical temperature is approached. This follows from the fact that the smaller the difference in the densities of vapor and liquid, the smaller the expansion that occurs during evaporation, and, therefore, the less work done during evaporation.

The existence of a critical temperature was first pointed out in 1860. Dmitry Ivanovich Mendeleev (1834-1907), Russian chemist who discovered the fundamental law of modern chemistry - the periodic law of chemical elements. Great achievements in the study of critical temperature belong to the English chemist Thomas Andrews, who carried out a detailed study of the behavior of carbon dioxide during an isothermal change in the volume it occupies. Andrews showed that at lower temperatures in a closed vessel the coexistence of carbon dioxide in liquid and gaseous states is possible; at higher temperatures such coexistence is impossible and the entire vessel is filled only with gas, no matter how much its volume is reduced.

After the discovery of the critical temperature, it became clear why gases such as oxygen or hydrogen could not be converted into liquid for a long time. Their critical temperature is very low (Table 23). To turn these gases into liquid, they must be cooled below a critical temperature. Without this, all attempts to liquefy them are doomed to failure.