Transitions between states of aggregation. Change in aggregative states of a substance Transition of a substance from gas to solid

The transition of a substance from a gaseous state to a liquid state is called condensation. Under certain conditions, substances can change from one state to another. Transition from one state of aggregation to another. The reverse process of sublimation (the transition of a substance from a gas to a solid state) is called desublimation. There are some substances that go from a solid state directly to a gaseous state, bypassing the liquid phase. This process is called sublimation or sublimation.

Evaporation can occur at any temperature. The transition of liquid into vapor, which occurs throughout the entire volume of the body, is called boiling, and the temperature at which the liquid boils is called the boiling point. If other parameters of the external environment (in particular, pressure) remain constant, then the body temperature does not change during the process of melting (crystallization) and boiling.

2. Liquid state

There are other states of aggregation, for example, the Bose-Einstein condensate. A distinctive feature is the absence of a sharp boundary of the transition to the plasma state. Definitions of states of aggregation are not always strict. Solids are divided into crystalline and amorphous. Crystals are characterized by spatial periodicity in the arrangement of equilibrium positions of atoms, which is achieved by the presence of long-range order and is called a crystal lattice.

1. Solid state

According to classical concepts, the stable state (with a minimum of potential energy) of a solid is crystalline. A special case of the amorphous state is the glassy state. The liquid state is usually considered intermediate between a solid and a gas. The shape of liquid bodies can be determined entirely or partly by the fact that their surface behaves like an elastic membrane.

As a rule, a substance in the liquid state has only one modification. Like gas, liquids are also mostly isotropic. However, there are liquids with anisotropic properties - liquid crystals. In addition to the isotropic, so-called normal phase, these substances, mesogens, have one or more ordered thermodynamic phases, which are called mesophases.

3. Gaseous state

Molecules in a gas can move freely and quickly. The gaseous state is characterized by the fact that it does not retain either shape or volume. Gas fills all available space and penetrates into any nooks and crannies. This is a state characteristic of substances with low density.

From a microscopic point of view, a gas is a state of matter in which its individual molecules interact weakly and move chaotically. Like liquids, gases have fluidity and resist deformation. Unlike liquids, gases do not have a fixed volume and do not form a free surface, but tend to fill the entire available volume (for example, a vessel). Some substances do not have a gaseous state.

The fourth state of matter is often called plasma. Liquid crystals simultaneously have the properties of both liquids (fluidity) and crystals (anisotropy). Structurally, liquid crystals are viscous liquids consisting of elongated or disk-shaped molecules, ordered in a certain way throughout the entire volume of this liquid.

Evaporation and condensation

In turn, nematics are divided into nematic and cholesteric liquid crystals. Since helium atoms are bosons, quantum mechanics allows an arbitrary number of particles to be in one state.

Since the energy of states is discrete, an atom can receive not any energy, but only one that is equal to the energy gap between adjacent energy levels. But at low temperatures, the collision energy may be less than this value, as a result of which energy dissipation simply will not occur.

It is also believed that in the evolution of the Universe, the state of glasma preceded the quark-gluon plasma, which existed in the first millionths of a second immediately after the Big Bang. During deep cooling, some (not all) substances transform into a superconducting or superfluid state. A fundamentally different state of matter, consisting only of neutrons.

The diffusion coefficient in this case has a value intermediate between liquid and gas. Substances in a supercritical state can be used as substitutes for organic solvents in laboratory and industrial processes.

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As a result, matter in the neutron state consists entirely of neutrons and has a density on the order of nuclear. MeV and higher) in the neutron state, various mesons begin to be born and annihilate. With a further increase in temperature, deconfinement occurs, and the substance passes into the state of quark-gluon plasma. It no longer consists of hadrons, but of constantly being born and disappearing quarks and gluons.

What is a “triple point” and how to determine its coordinates? Experiments show that for each substance there are conditions (pressure and temperature) under which vapor, liquid and crystal can coexist simultaneously for an arbitrarily long time. For example, if you place water with floating ice in a closed vessel at zero degrees, then both water and ice will evaporate into the free space. However, at a vapor pressure of 0.006 atm. (this is their “own” pressure, without taking into account the pressure created by the air) and a temperature of 0.01 ° C, the increase in the mass of steam will stop. From this moment on, ice, water and steam will retain their masses indefinitely. This is the triple point for water (left diagram). If water or steam is placed in the conditions of the left area, they will become ice. If you add a liquid or a solid to the “lower region,” you get steam. In the right area, water will condense and ice will melt.

A similar diagram can be constructed for any substance. The purpose of such diagrams is to answer the question: what state of matter will be stable at such and such pressure and such and such temperature. For example, the diagram on the right is for carbon dioxide. The triple point for this substance has a “pressure” coordinate of 5.11 atm, that is, significantly greater than normal atmospheric pressure. Therefore, under normal conditions (pressure 1 atm), we can only observe transitions “below the triple point,” that is, the independent transformation of a solid into a gas. At a pressure of 1 atm, this will occur at a temperature of –78 °C (see the dotted coordinate lines below the triple point).

We all live “near” the values of “normal conditions”, that is, primarily at a pressure close to one atmosphere. Therefore, if the atmospheric pressure is lower than the pressure corresponding to the triple point, when the body is heated, we will not see liquid - the solid will immediately turn into vapor. This is exactly how “dry ice” behaves, which is very convenient for ice cream sellers. Ice cream briquettes can be layered with pieces of “dry ice” and not be afraid that the ice cream will get wet. If the pressure corresponding to the triple point is less than atmospheric, then the substance is classified as “melting” - when the temperature rises, it first turns into liquid and then boils.

As you can see, the features of aggregate transformations of substances directly depend on how the current values of pressure and temperature relate to the coordinates of the “triple point” on the pressure-temperature diagram.

And in conclusion, let’s name substances known to you that always sublimate under normal conditions. This is iodine, graphite, “dry ice”. At pressures and temperatures different from normal, these substances can be observed in a liquid and even a boiling state.

Let's consider three states of aggregation: solid, liquid and gaseous and two transitions to them.

Phase transition "solid - liquid"

From the school physics course, four facts are known about this transition.

Fact one: the transition of a substance from a solid state (phase) to a liquid is called melting, and the reverse – crystallization.

Fact two: When melting, the system absorbs heat, and when solidifying, it releases heat.

Fact three: During the process of melting (crystallization), the temperature of the system remains constant until the entire system is melted. This temperature is called melting point.

Fact four: melting law: the amount of heat δQ required to melt a substance of mass dm is proportional to this mass:

(6.3.1)

The proportionality coefficient λ is a constant that depends only on the substance of the system and is called specific heat of fusion.

This law is also valid for crystallization, although with one difference: δ Q in this case, the heat generated by the system. Therefore, in generalized form the law can be written:

The change in entropy during this phase transition can be found simply by considering the process equilibrium.

This is a completely acceptable approximation, if we assume that the temperature difference between the system and the object that supplies heat to the system is not too great, much less than the melting point. Then we can use the thermodynamic meaning of entropy: from the point of view of thermodynamics, entropy is a function of the state of the system, the change of which d S in an elementary equilibrium process is equal to the ratio of the portion of heat δ Q, which the system receives in this process, to the system temperature T:

Substituting the expression for δQ here, we get:

From this formula it follows that during melting, entropy increases, and during crystallization it decreases. The physical meaning of this result is quite clear: the phase region of a molecule in a solid is much smaller than in a liquid, since in a solid, each molecule has access to only a small region of space between adjacent nodes of the crystal lattice, and in a liquid, molecules occupy the entire region of space. Therefore, at equal temperature, the entropy of a solid body is less than the entropy of a liquid. This means that a solid is a more ordered and less chaotic system than a liquid.

Phase transition "liquid - gas"

This transition has all the properties of a solid-liquid transition.

There are four facts that are familiar from school physics courses.

Fact one: the transition of a substance from a liquid to the gas phase is called evaporation, and the reverse transition is called condensation.

Fact two: When evaporating, the system absorbs heat, and when condensing, it loses.

Fact three: The processes of evaporation and condensation occur in a wide range of temperatures, but they are a phase transition only when the process involves the entire mass of the substance. This happens at a certain temperature T k, which is called boiling point. Each substance has its own boiling point. During the liquid-gas phase transition, the temperature remains constant and equal to the boiling point until the entire system passes from one phase to another.

Fact four: law of evaporation: the amount of heat δQ required to evaporate a substance of mass dm, and which is proportional to this mass:

Proportion factor r in this expression there is a constant depending on the substance of the system, called the specific heat of evaporation.

This law is also valid for condensation, although with one difference: δ Q in this case, the heat generated by the system. Therefore, the law of evaporation can be written in general form:

(6.3.3)

Where the plus sign refers to evaporation, and the minus sign refers to condensation.

The application of entropy in this process can be found simply by considering the process to be equilibrium. And again, this is a completely acceptable approximation, provided that the temperature difference between the system and the “supplier” of heat is small, i.e. much lower than boiling point. Then

6.3.4

From formula (6.3.4) it follows that during evaporation, entropy increases, and during condensation it decreases.

The physical meaning of this result is the difference in the phase region of the molecule in liquid and gas. Although in liquids and gases each molecule has access to the entire region of space occupied by the system, this region itself is significantly smaller for a liquid than for a gas. In a liquid, the attractive forces between molecules keep them at a certain distance from each other. Therefore, although each molecule has the opportunity to freely migrate through the region of space occupied by the liquid, it does not have the opportunity to “break away from the collective” of other molecules: as soon as it breaks away from one molecule, another is immediately attracted. Therefore, the volume of liquid depends on its quantity and is in no way related to the volume of the vessel.

Gas molecules behave differently. They have much more freedom, the average distance between them is such that the attractive forces are very small, and the molecules “notice each other” only during collisions. As a result, the gas always occupies the entire volume of the vessel.

Therefore, at equal temperatures, the phase region of gas molecules is significantly larger than the phase region of liquid molecules, and the entropy of the gas is greater than the entropy of the liquid. Gas, compared to liquid, is a much less ordered, more chaotic system.

It is important to know and understand how transitions between states of matter occur. We depict a diagram of such transitions in Figure 4.

5 - sublimation (sublimation) - transition from a solid to a gaseous state, bypassing the liquid;

6 - desublimation - transition from a gaseous state to a solid state, bypassing the liquid state.

B. 2 Ice melting and water freezing (crystallization)
If you place ice in a flask and start heating it with a burner, you will notice that its temperature will begin to rise until it reaches the melting point (0 o C). Then the melting process will begin, but the temperature of the ice will not increase, and only after the melting process of all the ice has completed, the temperature of the resulting water will begin to increase.

Definition. Melting- the process of transition from solid to liquid. This process occurs at a constant temperature.

The temperature at which a substance melts is called the melting point and is a measured value for many solids, and therefore a tabular value. For example, the melting point of ice is 0 o C, and the melting point of gold is 1100 o C.

The reverse process to melting - the process of crystallization - is also conveniently considered using the example of freezing water and turning it into ice. If you take a test tube with water and begin to cool it, then first there will be a decrease in the temperature of the water until it reaches 0 o C, and then it freezes at a constant temperature), and after complete freezing, further cooling of the formed ice.
If the described processes are considered from the point of view of the internal energy of the body, then during melting all the energy received by the body is spent on destroying the crystal lattice and weakening intermolecular bonds, thus, energy is spent not on changing temperature, but on changing the structure of the substance and the interaction of its particles. During the process of crystallization, energy exchange occurs in the opposite direction: the body gives off heat to the environment, and its internal energy decreases, which leads to a decrease in the mobility of particles, an increase in the interaction between them and solidification of the body.

Melting and crystallization graph

It is useful to be able to graphically depict the processes of melting and crystallization of a substance on a graph. The axes of the graph are: the abscissa axis is time, the ordinate axis is the temperature of the substance. As the substance under study, we will take ice at a negative temperature, i.e., ice that, upon receiving heat, will not immediately begin to melt, but will be heated to the melting temperature. Let us describe the areas on the graph that represent individual thermal processes:
Initial state - a: heating of ice to a melting point of 0 o C;
a - b: melting process at a constant temperature of 0 o C;
b - a point with a certain temperature: heating the water formed from ice to a certain temperature;
A point with a certain temperature - c: cooling of water to a freezing point of 0 o C;
c - d: the process of freezing water at a constant temperature of 0 o C;
d - final state: cooling of ice to a certain negative temperature.

Under normal conditions, any substance exists in one of three states - solid, liquid or gaseous ( cm. Aggregate states of matter). Each of these conditions corresponds to its own structure of bonds between molecules and/or atoms, characterized by a certain bond energy between them. To change this structure, either an influx of thermal energy from the outside is required (for example, during the melting of a solid substance), or an outflow of energy outward (for example, during crystallization).

Taking, for starters, a solid substance, we understand speculatively that the molecules/atoms in it are bound into some kind of rigid crystalline or amorphous structure - with slight heating they only begin to “shake” around their fixed position (the higher the temperature, the greater the amplitude of vibrations ). With further heating of the substance, the molecules loosen more and more until, finally, they break away from their “home” and go “free floating”. That's what it is melting or melting solid into liquid. The supply of energy necessary to melt a substance is called heat of fusion.

The graph of the change in temperature of a solid as it passes its melting point is in itself very interesting. Up to the melting point, as they heat up, the atoms/molecules swing around their fixed position more and more, and the arrival of each additional portion of thermal energy leads to an increase in the temperature of the solid. However, once a solid reaches its melting point, it remains at this temperature for some time, despite the continuing flow of heat, until it accumulates a sufficient amount of thermal energy to break rigid intermolecular bonds. That is, in the process phase transition a substance from a solid state to a liquid is absorbed by it without increasing the temperature, since all of it is spent on breaking intermolecular bonds. That’s why an ice cube in a cocktail, even in the hottest weather, remains icy in temperature until it’s all melted. At the same time, when melting, the ice cube takes away heat from the cocktail surrounding it (and thereby cools it to a pleasant temperature), and itself gains the energy that it requires to break intermolecular bonds and ultimately self-destruct.

The amount of heat required to melt or evaporate a unit volume of a solid or liquid is called, respectively, latent heat of fusion or latent heat of vaporization. And the quantities involved here are sometimes considerable. For example, to heat 1 kg of water from 0°C to 100°C requires “only” 420,000 joules (J) of thermal energy, and to turn this kilogram of water into 1 kg of steam with a temperature equal to the same 100°C , - as much as 2,260,000 J of energy.

After the solid mass has completely turned into a liquid, further heat will again lead to an increase in the temperature of the substance. In the liquid state, the molecules of a substance are still in close contact, but the rigid intermolecular bonds between them are broken, and the interaction forces holding the molecules together are several orders of magnitude weaker than in a solid, so the molecules begin to move quite freely relative to each other. Further supply of thermal energy brings the liquid to the phase boiling, and active evaporation or vaporization.

And, again, as was described in the case of melting or melting, for some time all the additional energy supplied is spent on breaking the liquid bonds between the molecules and releasing them into a gaseous state (at a constant boiling point). The energy spent on breaking these seemingly weak ties is the so-called. latent heat of vaporization - a considerable amount is also required (see example above).

All the same processes during the outflow of energy (cooling) of a substance occur in the reverse order. First, the gas cools as the temperature decreases, and this continues until it reaches condensation points- temperature at which it starts liquefaction, - and it is exactly equal to the evaporation (boiling) temperature of the corresponding liquid. During condensation, as the forces of mutual attraction between molecules begin to take precedence over the energy of thermal motion, the gas begins to turn into a liquid - “condense.” In this case, the so-called specific heat of condensation - it is exactly equal to the latent specific heat of evaporation, which has already been discussed. That is, how much energy you spent on evaporating a certain mass of liquid, exactly the same amount of energy the steam will give off in the form of heat when condensing back into the liquid.

The fact that the amount of heat released during condensation is very high is an easily verifiable fact: just raise your palm to the spout of a boiling kettle. In addition to the heat from the steam itself, your skin will also suffer from the heat released as a result of its condensation into liquid water.

As the liquid cools further to freezing points(whose temperature is equal to melting point), the process of releasing thermal energy outside will once again begin without lowering the temperature of the substance itself. This process is called crystallization, and it releases exactly the same amount of thermal energy as is taken from the environment during melting (the transition of a substance from a solid phase to a liquid).

There is another type of phase transition - from the solid state of a substance directly to the gaseous state (bypassing the liquid). This phase transformation is called sublimation, or sublimation. The most common example: damp laundry hung out to dry in the cold. The water in it first crystallizes into ice, and then - under the influence of direct sunlight - microscopic ice crystals simply evaporate, bypassing the liquid phase. Another example: at rock concerts, “dry ice” (frozen carbon dioxide CO 2) is used to create a smoke screen - it evaporates directly into the air, enveloping the performing musicians and also bypassing the liquid phase. Accordingly, it takes energy of sublimation.