Change of aggregative states of matter. Physical state The process of transition of a liquid into a solid state
Any change in the state of matter is associated with metamorphoses of temperature and pressure. One substance can be represented in the following states of aggregation: solid, liquid, gaseous.
Note that as the transition proceeds, no change in the composition of the substance is observed. The transition of a substance from a liquid to a solid state is accompanied only by a change in the forces of intermolecular interaction and the arrangement of molecules. The transformation from one state to another is called
Melting
This process involves turning into a liquid. For its implementation, elevated temperature is required.
For example, this state of matter can be observed in nature. Physics easily explains the process of melting snowflakes under the influence of spring rays. Small ice crystals that make up the snow begin to collapse after the air warms up to zero. Melting occurs gradually. First, the ice absorbs thermal energy. As the temperature changes, the ice completely transforms into liquid water.
It is accompanied by a significant increase in the speed of particle movement, thermal energy, and an increase in the value of internal energy.
After reaching the indicator, the structure of the solid substance breaks down. The molecules have greater freedom, they “jump”, occupying different positions. A molten substance has a greater energy reserve than in a solid state.
Curing temperature
The transition of a substance from a liquid to a solid state occurs at a certain temperature. If heat is removed from the body, it freezes (crystallizes).
The hardening temperature is considered one of the most important characteristics.
Crystallization
The transition of a substance from a liquid state to a solid state is called crystallization. When heat transfer to the liquid stops, a decrease in temperature to a certain value is observed. The phase transition of a substance from a liquid state to a solid is called crystallization in physics. When considering a substance that does not contain impurities, the melting point corresponds to the crystallization index.
Both processes occur gradually. The crystallization process is accompanied by a decrease in the molecules contained in the liquid. The attractive forces, due to which particles are held in strict order, inherent in solids, increase. Once the particles acquire an ordered arrangement, a crystal is formed.
They call the physical form of a substance present in a certain range of pressures and temperatures. It is characterized by quantitative properties that change in selected intervals:
- the ability of a substance to change shape and volume;
- absence (presence) of long-range or short-range order.
The crystallization process is associated with entropy, free energy, density, and other physical quantities.
In addition to liquids, solids, and gaseous forms, there is another state of aggregation - plasma. Gases can pass into it if the temperature increases at a constant pressure.
The boundaries between the various states of matter are not always strict. Physics has confirmed the existence of amorphous bodies capable of maintaining the structure of a liquid with low fluidity. have the ability to polarize electromagnetic radiation that passes through them.
Conclusion
In order to describe various states in physics, the definition of thermodynamic phase is used. Critical phenomena are states that describe the transformation of one phase into another. Solids are characterized by maintaining their average position over a long period of time. They will make slight oscillations (with minimal amplitude) around the equilibrium position. Crystals have a specific shape that will change when they become liquid. Information about boiling (melting) temperatures allows physicists to use transitions from one state of aggregation to another for practical purposes.
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.
Any body can be in different states of aggregation at a certain temperature and pressure - in solid, liquid, gaseous and plasma states.
For a transition from one state of aggregation to another, it occurs under the condition that the heating of the body from the outside occurs faster than its cooling. And vice versa, if the cooling of the body from the outside occurs faster than the heating of the body due to its internal energy.
When transitioning to another state of aggregation, the substance remains the same, the same molecules will remain, only their relative arrangement, speed of movement and forces of interaction with each other will change.
Those. a change in the internal energy of the particles of a body transfers it from one phase of the state to another. Moreover, this state can be maintained in a wide temperature range of the external environment.
When changing the state of aggregation, a certain amount of energy is needed. And during the transition process, energy is spent not on changing the body temperature, but on changing the internal energy of the body.
Let us display on the graph the dependence of body temperature T (at constant pressure) on the amount of heat Q supplied to the body during the transition from one state of aggregation to another.
Consider a body with mass m, which is in a solid state at a temperature T 1.
The body does not immediately transition from one state to another. First, energy is needed to change internal energy, and this takes time. The rate of transition depends on the mass of the body and its heat capacity.
Let's start heating the body. Using formulas you can write it like this:
Q = c⋅m⋅(T 2 -T 1)
The body must absorb so much heat in order to heat up from temperature T1 to T2.
Transition from solid to liquid
Further, at the critical temperature T2, which is different for each body, intermolecular bonds begin to break down and the body passes into another state of aggregation - liquid, i.e. intermolecular bonds weaken, molecules begin to move with greater amplitude, greater speed and greater kinetic energy. Therefore, the temperature of the same body in a liquid state is higher than in a solid state.
In order for the entire body to pass from a solid to a liquid state, it takes time to accumulate internal energy. At this time, all the energy goes not to heating the body, but to the destruction of old intermolecular bonds and the creation of new ones. Amount of energy needed:
λ - specific heat of melting and crystallization of a substance in J/kg, different for each substance.
After the entire body has passed into a liquid state, this liquid again begins to heat up according to the formula: Q = c⋅m⋅(T-T 2); [J].
Transition of a body from liquid to gaseous state
When a new critical temperature T 3 is reached, a new process of transition from liquid to vapor begins. To move further from liquid to vapor, you need to expend energy:
r is the specific heat of gas formation and condensation of a substance in J/kg, different for each substance.
Note that a transition from the solid state to the gaseous state is possible, bypassing the liquid phase. This process is called sublimation, and its inverse process is desublimation.
Transition of a body from a gaseous state to a plasma state
Plasma- a partially or fully ionized gas in which the densities of positive and negative charges are almost equal.
Plasma usually occurs at high temperatures, from several thousand °C and above. Based on the method of formation, two types of plasma are distinguished: thermal, which occurs when gas is heated to high temperatures, and gaseous, which is formed during electrical discharges in a gaseous environment.
This process is very complex and has a simple description, and it is not achievable for us in everyday conditions. Therefore, we will not dwell on this issue in detail.
All matter can exist in one of four forms. Each of them is a specific state of aggregation of a substance. In the nature of the Earth, only one is represented in three of them at once. This is water. It is easy to see both evaporated, and melted, and hardened. That is, steam, water and ice. Scientists have learned to change the aggregate states of matter. The biggest difficulty for them is only plasma. This condition requires special conditions.
What is it, what does it depend on and how is it characterized?
If a body has passed into a different state of matter, this does not mean that something else has appeared. The substance remains the same. If the liquid had water molecules, then the ice and steam would have the same molecules. Only their location, speed of movement and forces of interaction with each other will change.
When studying the topic “States of Aggregation (Grade 8),” only three of them are considered. These are liquid, gas and solid. Their manifestations depend on the physical conditions of the environment. The characteristics of these conditions are presented in the table.
| Name of the state of aggregation | solid | liquid | gas |
| Its properties | retains shape with volume | has a constant volume, takes the shape of a vessel | does not have constant volume and shape |
| Molecular arrangement | at the nodes of the crystal lattice | disorderly | chaotic |
| Distance between them | comparable to the size of molecules | approximately equal to the size of molecules | significantly larger than their size |
| How molecules move | oscillate around a lattice node | do not move from the point of equilibrium, but sometimes make large leaps | erratic with occasional collisions |
| How do they interact? | are strongly attracted | are strongly attracted to each other | do not attract, repulsive forces appear during impacts |
First state: solid
Its fundamental difference from others is that the molecules have a strictly defined place. When people talk about a solid state of aggregation, they most often mean crystals. Their lattice structure is symmetrical and strictly periodic. Therefore, it is always preserved, no matter how far the body spreads. The vibrational motion of the molecules of a substance is not enough to destroy this lattice.
But there are also amorphous bodies. They lack a strict structure in the arrangement of atoms. They could be anywhere. But this place is as stable as in the crystalline body. The difference between amorphous substances and crystalline substances is that they do not have a specific melting (solidification) temperature and are characterized by fluidity. Vivid examples of such substances: glass and plastic.
Second state: liquid
This state of matter is a cross between a solid and a gas. Therefore, it combines some properties from the first and second. Thus, the distance between particles and their interaction is similar to what was in the case of crystals. But the location and movement is closer to the gas. Therefore, the liquid does not retain its shape, but spreads throughout the vessel into which it is poured.
Third state: gas
For the science called “physics,” the state of aggregation in the form of gas is not in last place. After all, she studies the world around her, and the air in it is very widespread.
The peculiarities of this state are that there are practically no interaction forces between molecules. This explains their free movement. Due to which the gaseous substance fills the entire volume provided to it. Moreover, everything can be transferred to this state; you just need to increase the temperature by the required amount.
Fourth state: plasma
This state of aggregation of a substance is a gas that is completely or partially ionized. This means that the number of negatively and positively charged particles in it is almost the same. This situation occurs when gas is heated. Then there is a sharp acceleration of the process of thermal ionization. It consists in the fact that molecules are divided into atoms. The latter then turn into ions.
Within the Universe, this state is very common. Because it contains all the stars and the medium between them. It occurs extremely rarely within the boundaries of the Earth's surface. Apart from the ionosphere and solar wind, plasma is only possible during a thunderstorm. In lightning flashes, conditions are created in which atmospheric gases transform into the fourth state of matter.
But this does not mean that the plasma was not created in the laboratory. The first thing we managed to reproduce was a gas discharge. Plasma now fills fluorescent lamps and neon advertising.
How is the transition between states accomplished?
To do this, you need to create certain conditions: constant pressure and a specific temperature. In this case, a change in the aggregate state of a substance is accompanied by the release or absorption of energy. Moreover, this transition does not occur at lightning speed, but requires a certain amount of time. During this entire time, conditions must remain unchanged. The transition occurs with the simultaneous existence of a substance in two forms that maintain thermal equilibrium.
The first three states of matter can mutually transform into one another. There are direct processes and reverse ones. They have the following names:
- melting(solid to liquid) and crystallization, for example, melting ice and solidifying water;
- vaporization(from liquid to gaseous) and condensation, an example is the evaporation of water and its production from steam;
- sublimation(from solid to gas) and desublimation, for example, the evaporation of dry flavoring for the first of them and frosty patterns on the glass for the second.
Physics of melting and crystallization
If a solid is heated, then at a certain temperature, called melting point of a specific substance, a change in the state of aggregation will begin, which is called melting. This process involves the absorption of energy, which is called amount of heat and is designated by the letter Q. To calculate it you will need to know specific heat of fusion, which is denoted λ . And the formula takes on the following expression:
Q = λ * m, where m is the mass of the substance that is involved in melting.
If the reverse process occurs, that is, crystallization of the liquid, then the conditions are repeated. The only difference is that energy is released, and a minus sign appears in the formula.
Physics of vaporization and condensation
As the substance continues to be heated, it will gradually approach the temperature at which its intense evaporation begins. This process is called vaporization. It is again characterized by the absorption of energy. Only to calculate it you need to know specific heat of vaporization r. And the formula will be like this:
Q = r * m.
The reverse process or condensation occurs with the release of the same amount of heat. Therefore, a minus appears in the formula again.
The establishment of ideal order in the arrangement of atoms, i.e., the formation of a solid body, is prevented by thermal movements, the main feature of which, as we know, is chaoticity, disorder. Therefore, in order for a substance to be in a solid state, its temperature must be low enough - so low that the energy of thermal motion is less than the potential energy of interaction of atoms.
A body can only be a completely ideal crystal, in which all atoms are in equilibrium and have minimal energy, only at absolute zero. In reality, all substances become solid at much higher temperatures. The only exception is helium, which remains liquid even at absolute zero, but this is due to some quantum effects, which we will briefly discuss below.
A substance can pass into a solid state from either a liquid or a gaseous state. In both cases, such a transition is a transition from a state without symmetry to a state in which symmetry exists (this, in any case, refers to the long-range order that exists in crystals, but does not exist in either liquid or gaseous substances) . Therefore, the transition to the solid state must occur abruptly, that is, at a certain temperature, in contrast to the gas-liquid transition, which, as we know, can occur continuously.
Let us first consider the liquid-solid transformation. The process of formation of a solid when a liquid is cooled is the process of crystal formation (crystallization), (and it occurs at a certain temperature - the temperature of crystallization or solidification. Since during such a transformation the energy decreases, it is accompanied by the release of energy in the form of latent heat of crystallization. The reverse transformation is melting - also occurs abruptly at the same temperature and is accompanied by the absorption of energy in the form
that heat of fusion equal in value to the heat of crystallization.
This is clearly seen from the graph of the coolant temperature versus time shown in Fig. 179 (curve a). Section 1 of curve a gives the course of a monotonic decrease in the temperature of the liquid due to heat removal from it. Horizontal section 2 shows that at a certain temperature value, its decrease stops, despite the fact that heat removal continues. After some time, the temperature begins to drop again (section 3). The temperature corresponding to section 2 is the crystallization temperature. The heat released during crystallization compensates for the heat removal from the substance and therefore the temperature decrease temporarily stops. After the end of the crystallization process, the temperature of the now solid body begins to decrease again.
This progression of the temperature decrease graph is typical for crystalline bodies. When cooling liquids that do not crystallize (amorphous substances), latent heat is not released and the cooling graph is a monotonic curve without stopping cooling.
During the reverse process of transition of a substance from a solid to a liquid state (melting), a stop in the temperature increase is also observed on the heating curve, due to the absorption of the latent heat of melting - the heat due to which the crystal lattice is destroyed (curve in Fig. 179).
For crystallization to begin, the presence of a center or centers of crystallization is necessary. Such centers could be random accumulations of liquid particles stuck to each other, to which more and more new particles could join until all the liquid turns into a solid. However, the formation of such accumulations in the liquid itself is hampered by thermal movements, which destroy them even before they have time to acquire any noticeable size. Crystallization is greatly facilitated if sufficiently large solid particles in the form of dust particles and bodies are present in the liquid from the very beginning, which become centers of crystallization.
The formation of crystallization centers in the liquid itself is facilitated, of course, with a decrease in temperature. Therefore, the crystallization of a pure liquid, devoid of foreign formations,
usually begins at a temperature slightly lower than the true crystallization temperature. Under ordinary conditions, there are many crystallization centers in a crystallizing liquid, so that many crystals are formed in the liquid, accreting together, and the solidified substance turns out to be polycrystalline.
Only under special conditions, which are usually difficult to ensure, can a single crystal be obtained - a single crystal growing from a single crystallization center. If the same conditions for the accumulation of particles are provided for all directions, then the crystal turns out to be correctly faceted in accordance with its symmetry properties.
The liquid-solid transition, as well as the reverse transformation, is a phase transition, since the liquid and solid states can be considered as two phases of a substance. Both phases at the crystallization (melting) temperature can come into contact with each other, being in equilibrium (ice, for example, can float in water without melting), just as a liquid and its saturated vapor can be in equilibrium.
Just as the boiling point depends on pressure, the crystallization temperature (and its equal melting point) also depends on pressure, usually increasing with increasing pressure. It grows because external pressure brings the atoms closer together, and to destroy the crystal lattice during melting, the atoms need to be moved away from each other: at higher pressure, this requires greater energy of thermal motion, i.e., a higher temperature.
In Fig. 180 shows a curve of the melting temperature (crystallization) versus pressure. A solid curve divides the entire region into two parts. The area to the left of the curve corresponds to the solid state, and the area to the right of the curve corresponds to the liquid state. Any point lying on the melting curve itself corresponds to the equilibrium of the solid and liquid phases: at these pressures and temperatures, the substance in the liquid and solid states is in equilibrium, in contact with each other, and the liquid does not harden and the solid does not melt.
Dotted line in Fig. 180 shows the melting curve for those few substances (bismuth, antimony, ice, germanium) whose volume does not decrease during solidification, but increases. Such
substances, naturally, the melting point decreases with increasing pressure.
The change in melting temperature is related to the change in pressure by the Clapeyron-Clausius relation:
Here is the melting temperature (crystallization), and are, respectively, the molar volumes of the liquid and solid phases and the molar heat of fusion.
This formula is also valid for other phase transitions. In particular, for the case of evaporation and condensation, the Clapeyron-Clausius formula was derived in Chap. VII [see (105.6)].
From the Clapeyron-Clausius formula it is clear that the sign of the change in melting temperature with a change in pressure is determined by which of the two values, or greater. The steepness of the curve also depends on the value of the latent heat of transition; the less, the less the melting temperature changes with pressure. In table 20 shows the values of the specific (i.e., per unit mass) heat of fusion for some substances.
Table 20 (see scan) Specific heat of fusion for some substances
The Clapeyron-Clausius equation can also be written in this form:
This equation shows how the pressure under which both equilibrium phases are changes with temperature changes.
A solid can be formed not only by crystallization of a liquid, but also by condensation of gas (steam) into a crystal, bypassing the liquid phase. In this case, the latent heat of transition is also released, which, however, is always greater than the latent heat of fusion. After all, the formation of a solid at a certain temperature and pressure can occur either directly from a gaseous state or through preliminary liquefaction. In both
cases, the initial and final states are the same. This means that the energy difference between these states is also the same. Meanwhile, in the second case, firstly, the latent heat of condensation is released during the transition from the gaseous to the liquid state and, secondly, the latent heat of crystallization during the transition from the liquid to the solid state. It follows that the latent heat during the direct formation of a solid from the gaseous phase must be equal to the sum of the heat of condensation and crystallization from the liquid. This only applies to heats measured at the melting point. At lower temperatures, the heat of condensation from the gas increases.
The reverse process of evaporation of a solid is usually called sublimation or sublimation. Evaporating particles of a solid form steam above it in exactly the same way as occurs during the evaporation of a liquid. At certain pressure and temperature, vapor and solid can be in equilibrium. Steam in equilibrium with a solid is also called saturated steam. As in the case of a liquid, the saturated vapor pressure over a solid depends on temperature, decreasing rapidly with decreasing temperature, so that for many solids at ordinary temperatures the saturated vapor pressure is negligible.
In Fig. 181 shows the curve of the dependence of saturated vapor pressure on temperature. This curve is the equilibrium line between the solid and gaseous phases. The area to the left of the curve corresponds to the solid state, to the right of it - to the gaseous state. Sublimation, like melting, is associated with the destruction of the lattice and requires the expenditure of energy necessary for this. This energy manifests itself as the latent heat of sublimation (sublimation), equal, of course, to the latent heat of condensation. The heat of sublimation is therefore equal to the sum of the heats of fusion and vaporization.
