S.A. Vaporization
All gases are vapors of any substance, therefore there is no fundamental difference between the concepts of gas and vapor. Water vapor is a phenomenon. real gas and is widely used in various industries. This is explained by the ubiquity of water, its cheapness and harmlessness to human health. Water vapor is produced by the evaporation of water when heat is supplied to it.
Vaporization called the process of changing liquid into vapor.
Evaporation called vaporization that occurs only from the surface of the liquid and at any temperature. The intensity of evaporation depends on the nature of the liquid and temperature.
Boiling called vaporization throughout the entire mass of liquid.
The process of converting steam into liquid, which occurs when heat is removed from it and is a process reverse to vaporization, called. condensation. This process, as well as vaporization, occurs at a constant temperature.
Sublimation or sublimation called the process of a substance changing from a solid state directly to a vapor.
The process is the reverse of the sublimation process, i.e. the process of transition of steam directly into a solid state, called. desublimation.
Saturated steam. When a liquid evaporates into a limited volume, the reverse process also occurs simultaneously, i.e. liquefaction phenomenon. As steam evaporates and fills the space above the liquid, the intensity of evaporation decreases and the intensity of the reverse process increases. At some point, when the rate of condensation becomes equal to the rate of evaporation, dynamic equilibrium occurs in the system. In this state, the number of molecules flying out of the liquid will be equal to the number of molecules returning back into it. Consequently, in the vapor space at this equilibrium state there will be a maximum number of molecules. Steam in this state has a maximum density and is called. rich. By saturated we mean steam that is in equilibrium with the liquid from which it is formed. Saturated steam has a temperature that is a function of its pressure, equal to the pressure of the medium in which the boiling process occurs. When the volume of saturated vapor increases at a constant temperature, a certain amount of liquid transforms into vapor, and when the volume decreases at a constant temperature, the vapor transforms into liquid, but in both the first and second cases, the vapor pressure remains constant.
Dry saturated steam obtained when all the liquid evaporates. The volume and temperature of dry steam are functions of pressure. As a result, the state of dry steam is determined by one parameter, for example, pressure or temperature.
Wet saturated steam, resulting from incomplete evaporation of a liquid, phenomenon. a mixture of steam with tiny droplets of liquid, distributed evenly throughout its entire mass and suspended in it.
The mass fraction of dry steam in wet steam is called. degree of dryness or mass vapor content and is denoted by x. The mass fraction of liquid in wet vapor is called. degree of humidity and is denoted by y. Obviously y=1-x. The degree of dryness and the degree of humidity are expressed either as fractions of a unit or as a percentage.
For dry steam x=1, and for water x=0. During the process of steam formation, the degree of steam dryness gradually increases from zero to one.
When heat is imparted to dry steam at constant pressure, its temperature will increase. The steam produced in this process is called. overheated.
Since the specific volume of superheated steam is greater than the specific volume of saturated steam (since р=const, tper>tн), then the density of superheated steam is less than the density of saturated steam. Therefore, superheated steam is unsaturated. In terms of its physical properties, superheated steam approaches ideal gases.
10.3. R, v– water vapor diagram
Let's consider the features of the vaporization process. Let there be 1 kg of water in a cylinder at a temperature of 0 C, on the surface of which a pressure p is applied using a piston. The volume of water located under the piston is equal to the specific volume at 0 C, denoted by ( = 0.001 m / kg). For simplicity, we assume that water is a phenomenon. a practically incompressible liquid and has the highest density at 0 C, and not at 4 C (more precisely 3.98 C). When the cylinder is heated and heat is transferred to the water, its temperature will rise, the volume will increase, and when t = t n, corresponding to p = p 1, is reached, the water will boil and steam formation will begin.
All changes in the state of liquid and vapor will be noted in p, v coordinates (Fig. 10.1).
The process of formation of superheated steam at p=const consists of three sequentially carried out physical processes:
1. Heating the liquid to temperature tn;
2. Vaporization at t n =const;
3. Overheating of steam, accompanied by an increase in temperature.
When p=p 1 these processes in p, v– the diagram corresponds to segments a-a, a-a, a-d. In the interval between points a and a, the temperature will be constant and equal to tn1 and the steam will be wet, and closer to t.a its degree of dryness will be less (x = 0), and in t.a, corresponding to the state of dry steam, x = 1. If the vaporization process occurs at a higher pressure (p 2 >p 1), then the volume of water will practically remain the same. The volume v corresponding to boiling water will increase slightly (), because t n2 >t n1, and volume, since the process of vaporization at higher pressure and high temperature occurs more intensely. Consequently, as pressure increases, the volume difference (segment ) increases, and the volume difference (segment ) decreases. A similar picture will occur when the vaporization process occurs at higher pressure (p 3 >p 2 ; ; , because t n3 >t n2).
If in Fig. 10.1 we connect points with one and two strokes lying on the isobars
different pressures, we get lines ; ,
each of which has a very specific meaning. For example, line a-b-c expresses the dependence of the specific volume of water at 0 C on pressure. It is almost parallel to the ordinate axis, because Water is a practically incompressible liquid. The line shows the dependence of the specific volume of boiling water on pressure. This line is called lower boundary curve. In p, v– diagram, this curve separates the region of water from the region of saturated vapor. The line shows the dependence of the specific volume of dry steam on pressure and called. upper boundary curve. It separates the region of saturated steam from the region of superheated (unsaturated) steam.
The meeting point of boundary curves is called. critical point K. This point corresponds to a certain limiting critical state of the substance, when there is no difference between liquid and vapor. At this point there is no section of the vaporization process. The parameters of the substance in this state are called. critical. For example, for water: pk=22.1145 MPa; Tk=647.266 K; Vк=0.003147 m/kg.
Critical temperature maximum saturated steam temperature. At temperatures above the critical temperature, only superheated vapors and gases can exist. The concept of critical temperature was first given in 1860 by D.I. Mendeleev. He defined it as the temperature above which a gas cannot be converted into a liquid, no matter how high pressure is applied to it.
However, the vaporization process does not always occur as shown in Fig. 10.1. if the water is cleared of mechanical impurities and gases dissolved in it, vaporization can begin at a temperature above Tn (sometimes by 15-20 K) due to the absence of vaporization centers. This water is called overheated. On the other hand, with rapid isobaric cooling of superheated steam, its condensation may not begin at Tn. and at a slightly lower temperature. This pair is called hypothermic or oversaturated. When deciding what state of aggregation substances (steam or water) can be in at given p and T p and v or T and V you should always keep the following in mind. When p=const for superheated steam and T d >T n (see Fig. 10.1); for water, vice versa and T<Т н; при Т=const для перегретого пара и р е <р н; для воды и р n >rn. Knowing these relationships and using the tables for saturated steam, you can always determine in which of the three regions 1, 2 or 3 (see Fig. 10.2) the working fluid with the given parameters is located, i.e. whether it is liquid (region 1), saturated (region 2) or superheated (region 3) vapor.
For the supercritical region, the critical isotherm (dash-dotted curve) is conventionally taken as the probable water-steam boundary. In this case, to the left and to the right of this isotherm, the substance is in a single-phase homogeneous state, possessing, for example, in point y the properties of a liquid, and in point z – the properties of a vapor.
What is evaporation
Definition
Evaporation is the process of vaporization that occurs from the free surface of a liquid.
Evaporation occurs at any temperature and occurs more intensely with increasing temperature. As a result of evaporation, the liquid cools, since evaporation can be explained by the fact that molecules with the greatest kinetic energy fly out from the surface layer of the liquid, overcoming the attractive forces of neighboring molecules. The rate of evaporation depends on the external pressure and the movement of the gaseous phase above the free surface of the liquid. With increasing temperature, the density, and therefore the saturated vapor pressure above the liquid, increases. As the vapor density increases, the surface tension of the liquid decreases, therefore, the latent heat of vaporization decreases with increasing temperature. At the critical temperature ($(\T)_k$), the density of saturated vapors is equal to the density of the liquid, the difference between these phases of the substance disappears. It turns out that at the critical temperature, surface tension and latent heat of vaporization are equal to zero. Steam, strictly speaking, is not a gas. For vapors close to saturation, the pressure changes slightly depending on the volume. Gas laws can be approximately applied to unsaturated vapors.
What is boiling
Definition
The process of intense evaporation of a liquid not only from its free surface, but throughout the entire volume of the liquid into the vapor bubbles formed in the process is called boiling.
The pressure p inside the vapor bubble is determined in accordance with the following expression:
where $p_0$ is the external pressure, $\rho gh$ is the pressure of the layers of liquid that are located above, $p_(R\ )=\frac(2\sigma)(r)$ is the additional pressure that is caused by the curvature of the bubble , $r$ is the radius of the bubble, $h$ is the distance from the center of the bubble to the surface of the liquid, $\rho $ is the density of the liquid, $ \sigma $ is the surface tension of the liquid.
Boiling begins when the pressure (elasticity) of the saturated vapor inside the bubble ($p_p$) is greater than the pressure on the right side of formula (1). If a liquid has centers of vaporization, then boiling of the liquid begins at lower temperatures. If $\rho gh\ll p_0$, then we can assume that boiling begins at $p_p\approx p_0$. The temperature of a liquid at which its saturated vapor pressure is equal to external pressure is called the boiling temperature (point) ($(\T)_k$). Strictly speaking, boiling at different levels of a liquid occurs at different temperatures; there is no one specific temperature. Saturated steam, which is located above the surface of a boiling liquid, has a certain temperature. Its temperature does not depend on how boiling occurs in the depths of the liquid, and is determined only by external pressure. It is the temperature of such steam that is meant when talking about the boiling point.
If boiling occurs at constant pressure ($p_0$), then the boiling temperature is constant. The heat supplied to the system in this case is spent only on steam formation.
What is condensation
Definition
The reverse process of evaporation is called condensation.
Condensation releases heat. Liquid boiling and vapor condensation are first-order phase transitions. Let us recall that a first-order phase transition is a transition that is accompanied by an abrupt change in the internal energy and density of a substance. During first-order phase transitions, which include evaporation and condensation, the thermodynamic potential (Ф) of the system does not change.
The amount of heat that must be consumed during the vaporization of a unit mass of liquid at a temperature equal to $(\T)_k$ is called the specific heat of vaporization (or latent heat of boiling) ($r_k$). The specific heat of vaporization can be found from the Clayperon-Clausius equation:
where $v_p,v_j$ are the specific volumes of vapor and liquid at the boiling point $T_k$. Accordingly, the dependence of the boiling point on pressure during the evaporation process is defined as:
\[\frac(dT_k)(dp)=\frac(\left(v_p-v_j\right)T_k)(r_k)\ \left(3\right).\]
Since $v_p>v_j$ and $r_k>0$, then $\frac(dT_k)(dp)>0.\ $ Figure 1 shows the phase equilibrium curve of the vaporization process. It ends at the critical point K. The consequence of the break in the evaporation curve at point K is the continuity of the liquid and gaseous state of the substance. The boiling point increases with increasing pressure.
Example 1
Assignment: Determine the molar heat of evaporation of a liquid at temperature T and pressure p of saturated vapor, if the liquid and its vapor obey the van der Waals equation. The coefficients in the van der Waals equation are a and b, $T\ll T_k$.
The heat of vaporization, based on the first law of thermodynamics, can be calculated as:
where $U_j-$ internal energy of liquid, $U_p$ internal energy of vapor, $V_j,V_p$ volumes of liquid and vapor respectively, $p\left(V_p-V_j\right)-\work\performed\during\evaporation$ against external pressure forces p. The difference in internal energies according to the van der Waals equation is equal to:
We use the van der Waals equation for one mole of a substance:
\[\left(p+\frac(a)(V^2)\right)\left(V-b\right)=RT\to p=\frac(RT)(\left(V-b\right))-\frac( a)(V^2)\ \left(1.3\right).\]
We get:
Answer: The molar heat of evaporation of a liquid under given conditions is equal to: $r_p=V_p\left\(\frac(RT)(V_p-b)-\frac(2a)(V^2_p)\right\)-V_j\ \left\ (\frac(RT)(V_j-b)-\frac(2a)(V^2_j)\right\)$.
Example 2
Task: Two kilograms of water were taken at a temperature of 00C at atmospheric pressure, heated and completely turned into steam. Find the change in entropy if the process is considered reversible.
\[\triangle S=\int(\frac(\delta Q)(T))=\triangle S_(nagr)+\triangle S_p\ (2.1),\]
where $\triangle S_(nagr)$ is the change in entropy when water is heated from zero Celsius to boiling point, that is, from $T_1=273\ K\ to\ T_2=373\ K$. $\triangle S_p-\ $change in entropy during vaporization. Let's find the change in entropy when water is heated:
\[\triangle S_(nagr)=\int\limits^(T_2)_(T_1)(\frac(\delta Q)(T))=\int\limits^(T_2)_(T_1)(\frac( cmdT)(T))=cm\int\limits^(T_2)_(T_1)(\frac(dT)(T))=cmln\left(\frac(T_2)(T_1)\right)\left(2.2 \right),\]
where $c$ specific heat capacity of water is equal to $c=4.2\ (\cdot 10)^3\frac(J)(kgK)$
The process of vaporization at the boiling point occurs without a change in temperature, so the expression for the change in entropy in this process will be:
\[\triangle S_p=\frac(1)(T_2)\int(\delta Q)=\frac(1)(T_2)\triangle Q=\frac(r_pm)(T_2)\left(2.3\right), \]
where $r_p$ is the specific heat of vaporization from reference materials and is equal to $r_p=22.6\cdot (10)^5\frac(J)(kg)$ for water.
The final expression for the change in entropy is:
\[\triangle S=cmln\left(\frac(T_2)(T_1)\right)+\frac(r_pm)(T_2)\ \left(2.4\right).\]
All data has been converted to SI, let’s carry out the calculation:
\[\triangle S=4.2\ (\cdot 10)^3\cdot 2(ln \left(\frac(373)(273)\right)\ )+\frac(22.6\cdot (10) ^5\cdot 2)(373)=14.6(\cdot 10)^3\frac(J)(K).\]
Answer: The change in entropy in a given process is equal to $1.46(\cdot 10)^4\frac(J)(K)$.
Topics of the Unified State Examination codifier: change in aggregate states of matter, melting and crystallization, evaporation and condensation, boiling of liquid, change in energy in phase transitions.
Ice, water and water vapor are examples of the three states of aggregation substances: solid, liquid and gaseous. What exact state of aggregation a given substance is in depends on its temperature and other external conditions in which it is located.
When external conditions change (for example, if the internal energy of a body increases or decreases as a result of heating or cooling), phase transitions can occur - changes in the aggregate states of the body's substance. We will be interested in the following phase transitions.
Melting(solid-liquid) and crystallization(liquid-solid).
Vaporization(liquid vapor) and condensation(steam liquid).
Melting and crystallization
Most solids are crystalline, i.e. have crystal lattice- a strictly defined, periodically repeated arrangement of its particles in space.
Particles (atoms or molecules) of a crystalline solid undergo thermal vibrations near fixed equilibrium positions - nodes crystal lattice.
For example, the nodes of the crystal lattice of table salt are the vertices of cubic cells of “three-dimensional checkered paper” (see Fig. 1, in which larger balls represent chlorine atoms (image from en.wikipedia.org.)); If you let the water from the salt solution evaporate, the remaining salt will be a pile of small cubes.
Rice. 1. Crystal lattice
Melting called the transformation of a crystalline solid into a liquid. Any body can be melted - to do this you need to heat it to melting point, which depends only on the substance of the body, but not on its shape or size. The melting point of a given substance can be determined from tables.
On the contrary, if you cool a liquid, sooner or later it will turn into a solid state. The transformation of a liquid into a crystalline solid is called crystallization or hardening. Thus, melting and crystallization are mutually inverse processes.
The temperature at which liquid crystallizes is called crystallization temperature. It turns out that the crystallization temperature is equal to the melting temperature: at a given temperature, both processes can occur. So, when ice melts, water crystallizes; What exactly occurs in each specific case - depends on external conditions (for example, whether heat is supplied to the substance or removed from it).
How do melting and crystallization occur? What is their mechanism? To understand the essence of these processes, let us consider graphs of the dependence of body temperature on time during its heating and cooling - the so-called melting and crystallization graphs.
Melt graph
Let's start with the melting graph (Fig. 2). Let at the initial moment of time (point on the graph) the body be crystalline and have a certain temperature.
Rice. 2. Melting graph
Then heat begins to be supplied to the body (say, the body is placed in a melting furnace), and the body temperature rises to a value - the melting temperature of the given substance. This is a section of the graph.
At the site the body receives the amount of heat
where is the specific heat capacity of the solid substance, and is the mass of the body.
When the melting temperature is reached (at point ) the situation changes qualitatively. Despite the fact that heat continues to be supplied, body temperature remains unchanged. Happening at the site melting body - its gradual transition from solid to liquid. Inside the area we have a mixture of solid and liquid, and the closer to point , the less solid remains and the more liquid appears. Finally, at a point there was nothing left of the original solid body: it completely turned into a liquid.
The area corresponds to further heating of the liquid (or, as they say, melt). In this area, the liquid absorbs an amount of heat
where is the specific heat capacity of the liquid.
But what we are most interested in now is the phase transition area. Why doesn't the temperature of the mixture change in this area? The heat is coming!
Let's go back to the beginning of the heating process. An increase in the temperature of a solid body in an area is the result of an increase in the intensity of vibrations of its particles at the nodes of the crystal lattice: the supplied heat goes to increase kinetic energy of the particles of the body (in fact, some part of the supplied heat is spent on doing work to increase the average distances between particles - as we know, bodies expand when heated. However, this part is so small that it can be ignored.).
The crystal lattice becomes looser more and more, and at the melting temperature the range of vibrations reaches the limiting value at which the attractive forces between the particles are still able to ensure their ordered arrangement relative to each other. The solid body begins to “crack at the seams”, and further heating destroys the crystal lattice - this is how melting begins in the area.
From this moment, all the heat supplied is used to perform work on breaking the bonds that hold the particles in the nodes of the crystal lattice, i.e. to increase potential particle energy. The kinetic energy of the particles remains the same, so the body temperature does not change. At a point, the crystalline structure disappears completely, there is nothing left to destroy, and the supplied heat again goes to increase the kinetic energy of the particles - to heat the melt.
Specific heat of fusion
So, to transform a solid into a liquid, it is not enough to bring it to the melting point. It is necessary to additionally (already at the melting temperature) provide the body with a certain amount of heat for the complete destruction of the crystal lattice (i.e. to pass through the section).
This amount of heat goes to increase the potential energy of particle interaction. Consequently, the internal energy of the melt at a point is greater than the internal energy of the solid at a point by an amount.
Experience shows that the value is directly proportional to body weight:
The proportionality coefficient does not depend on the shape and size of the body and is a characteristic of the substance. It is called specific heat of fusion of a substance. The specific heat of fusion of a given substance can be found in the tables.
The specific heat of fusion is numerically equal to the amount of heat required to transform one kilogram of a given crystalline substance brought to the melting point into liquid.
Thus, the specific heat of melting of ice is equal to kJ/kg, of lead - kJ/kg. We see that it takes almost twice as much energy to destroy the ice crystal lattice! Ice is a substance with a high specific heat of fusion and therefore does not melt immediately in the spring (nature took its own measures: if ice had the same specific heat of fusion as lead, the entire mass of ice and snow would melt with the first thaw, flooding everything around).
Crystallization graph
Now let's move on to consider crystallization- a process reverse to melting. We start from the point of the previous drawing. Let us assume that at the point the heating of the melt has stopped (the stove has been turned off and the melt has been exposed to air). Further changes in the melt temperature are shown in Fig. (3) .
Rice. 3. Crystallization graph
The liquid cools (section) until its temperature reaches the crystallization temperature, which coincides with the melting point.
From this moment on, the temperature of the melt stops changing, although heat still escapes from it into the environment. Happening at the site crystallization melt - its gradual transition to a solid state. Inside the area we again have a mixture of solid and liquid phases, and the closer to the point, the more solid becomes and the less liquid becomes. Finally, at the point there is no liquid left at all - it has completely crystallized.
The next section corresponds to the further cooling of the solid body resulting from crystallization.
We are again interested in the phase transition section: why does the temperature remain unchanged despite the loss of heat?
Let's return to the point again. After the heat supply is stopped, the temperature of the melt decreases, as its particles gradually lose kinetic energy as a result of collisions with environmental molecules and the emission of electromagnetic waves.
When the temperature of the melt drops to the crystallization temperature (point), its particles will slow down so much that the forces of attraction will be able to “unfold” them properly and give them a strictly defined mutual orientation in space. This will create conditions for the emergence of a crystal lattice, and it will actually begin to form due to the further release of energy from the melt into the surrounding space.
At the same time, a counter process of energy release will begin: when the particles take their places at the nodes of the crystal lattice, their potential energy sharply decreases, due to which their kinetic energy increases - the crystallizing liquid is a source of heat (you can often see birds sitting near the ice hole. They warm themselves there!) . The heat released during crystallization exactly compensates for the heat loss to the environment, and therefore the temperature in the area does not change.
At the point, the melt disappears, and along with the completion of crystallization, this internal “generator” of heat also disappears. Due to the continued dissipation of energy into the external environment, the decrease in temperature will resume, but only the formed solid body (section) will cool down.
As experience shows, during crystallization in the area, exactly the same the amount of heat that was absorbed during melting in the area.
Vaporization and condensation
Vaporization is the transition of a liquid into a gaseous state (in steam). There are two ways of vaporization: evaporation and boiling.
Evaporation called vaporization, which occurs at any temperature from free surface liquids. As you remember from the sheet “Saturated Steam”, the cause of evaporation is the departure from the liquid of the fastest molecules that are able to overcome the forces of intermolecular attraction. These molecules form vapor above the surface of the liquid.
Different liquids evaporate at different speeds: the greater the force of attraction of molecules to each other, the fewer molecules per unit time will be able to overcome them and fly out, and the lower the evaporation rate. Ether, acetone, and alcohol evaporate quickly (they are sometimes called volatile liquids), water evaporates more slowly, and oil and mercury evaporate much more slowly than water.
The rate of evaporation increases with increasing temperature (in hot weather, laundry will dry faster), since the average kinetic energy of liquid molecules increases, and thus the number of fast molecules capable of leaving its limits increases.
The rate of evaporation depends on the surface area of the liquid: the larger the area, the more molecules have access to the surface, and evaporation occurs faster (which is why when hanging laundry, it is carefully straightened out).
Simultaneously with evaporation, the reverse process is also observed: the vapor molecules, making random movements above the surface of the liquid, partially return back to the liquid. The transformation of vapor into liquid is called condensation.
Condensation slows down the evaporation of a liquid. So, in dry air, laundry will dry faster than in humid air. It will dry faster in the wind: the steam is carried away by the wind, and evaporation occurs more intensely
In some situations, the rate of condensation may be equal to the rate of evaporation. Then both processes compensate each other and dynamic equilibrium occurs: the liquid does not evaporate from a tightly sealed bottle for years, and in this case there is saturated steam.
We constantly observe the condensation of water vapor in the atmosphere in the form of clouds, rain and dew that falls in the morning; It is evaporation and condensation that ensure the water cycle in nature, supporting life on Earth.
Since evaporation is the departure of the fastest molecules from the liquid, during the evaporation process the average kinetic energy of liquid molecules decreases, i.e. the liquid cools down. You are well familiar with the feeling of coolness and sometimes even chilliness (especially in the wind) when you come out of the water: water, evaporating over the entire surface of the body, carries away heat, while the wind accelerates the evaporation process (it’s now clear why we blow on hot tea. By the way, It’s even better to draw air into yourself, since dry ambient air then comes to the surface of the tea, and not moist air from our lungs ;-)).
You can feel the same coolness if you run a piece of cotton wool soaked in a volatile solvent (say, acetone or nail polish remover) over your hand. In forty-degree heat, thanks to the increased evaporation of moisture through the pores of our body, we maintain our temperature at a normal level; Without this thermoregulatory mechanism, in such heat we would simply die.
On the contrary, during the process of condensation, the liquid heats up: when the vapor molecules return to the liquid, they are accelerated by attractive forces from nearby liquid molecules, as a result of which the average kinetic energy of the liquid molecules increases (compare this phenomenon with the release of energy during crystallization of a melt!).
Boiling
Boiling- this is the vaporization that occurs throughout the entire volume liquids.
Boiling is possible because a certain amount of air is always dissolved in a liquid, which gets there as a result of diffusion. When the liquid is heated, this air expands, air bubbles gradually increase in size and become visible to the naked eye (in a pan of water they settle on the bottom and walls). Inside the air bubbles there is saturated steam, the pressure of which, as you remember, increases rapidly with increasing temperature.
The larger the bubbles become, the greater the Archimedean force acts on them, and at a certain moment the bubbles begin to separate and float up. Rising upward, the bubbles enter less heated layers of the liquid; the vapor in them condenses, and the bubbles shrink again. The collapse of the bubbles causes the familiar noise that precedes the boiling of the kettle. Finally, over time, the entire liquid warms up evenly, the bubbles reach the surface and burst, throwing out air and steam - the noise is replaced by gurgling, the liquid boils.
The bubbles thus serve as “conductors” of vapor from inside the liquid to its surface. During boiling, along with normal evaporation, the liquid is converted into steam throughout the entire volume - evaporation into air bubbles, followed by the release of steam out. This is why boiling liquid evaporates very quickly: a kettle, from which the water would evaporate for many days, will boil away in half an hour.
Unlike evaporation, which occurs at any temperature, a liquid begins to boil only when it reaches boiling point- exactly the temperature at which air bubbles are able to float and reach the surface. At the boiling point, the saturated vapor pressure becomes equal to the external pressure on the liquid(in particular, atmospheric pressure). Accordingly, the greater the external pressure, the higher the temperature at which boiling will begin.
At normal atmospheric pressure (atm or Pa), the boiling point of water is . That's why the pressure of saturated water vapor at temperature is Pa. This fact must be known to solve problems - it is often considered known by default.
At the top of Elbrus, the atmospheric pressure is atm, and water there will boil at a temperature of . And under pressure atm, water will begin to boil only at .
The boiling point (at normal atmospheric pressure) is a strictly defined value for a given liquid (boiling points given in the tables of textbooks and reference books are the boiling points of chemically pure liquids. The presence of impurities in a liquid can change the boiling point. For example, tap water contains dissolved chlorine and some salts, so its boiling point at normal atmospheric pressure may differ slightly from ). So, alcohol boils at , ether - at , mercury - at . Please note: the more volatile a liquid is, the lower its boiling point. In the table of boiling points we also see that oxygen boils at. This means that at normal temperatures oxygen is a gas!
We know that if the kettle is removed from the heat, the boiling will immediately stop - the boiling process requires a continuous supply of heat. At the same time, the temperature of the water in the kettle stops changing after boiling, remaining the same all the time. Where does the supplied heat go?
The situation is similar to the melting process: heat is used to increase the potential energy of the molecules. In this case - to perform work to remove molecules at such distances that the forces of attraction will be unable to keep the molecules close to each other, and the liquid will turn into a gaseous state.
Boiling graph
Let's consider a graphical representation of the process of heating a liquid - the so-called boiling chart(Fig. 4).
Rice. 4. Boiling graph
The section precedes the onset of boiling. In the area, the liquid boils, its mass decreases. At this point the liquid boils away completely.
To pass the section, i.e. In order for a liquid brought to the boiling point to completely turn into steam, a certain amount of heat must be supplied to it. Experience shows that a given amount of heat is directly proportional to the mass of the liquid:
The proportionality factor is called specific heat of vaporization liquids (at boiling point). The specific heat of vaporization is numerically equal to the amount of heat that must be supplied to 1 kg of liquid taken at boiling point in order to completely convert it into steam.
So, at the specific heat of vaporization of water is equal to kJ/kg. It is interesting to compare it with the specific heat of melting of ice (kJ/kg) - the specific heat of vaporization is almost seven times greater! This is not surprising: after all, to melt ice, you only need to destroy the ordered arrangement of water molecules at the nodes of the crystal lattice; at the same time, the distances between the molecules remain approximately the same. But to turn water into steam, you need to do much more work to break all the bonds between molecules and remove the molecules to significant distances from each other.
Condensation graph
The process of steam condensation and subsequent cooling of the liquid looks on the graph symmetrically to the process of heating and boiling. Here's the relevant one condensation graph for the case of one hundred degree water vapor, which is most often encountered in problems (Fig. 5).
Rice. 5. Condensation graph
At the point we have water vapor at . There is condensation in the area; inside this area there is a mixture of steam and water at . At the point there is no more steam, there is only water at . The area is the cooling of this water.
Experience shows that during the condensation of a mass vapor (i.e., when passing through a section), exactly the same amount of heat is released that was spent on converting a liquid mass into vapor at a given temperature.
Let's compare the following amounts of heat for fun:
Which is released when water vapor condenses;
, which is released when the resulting 100-degree water cools to a temperature of, say, .
J;
J.
These numbers clearly show that a steam burn is much worse than a boiling water burn. When boiling water comes into contact with the skin, “only” is released (the boiling water cools down). But in case of a burn with steam, an order of magnitude greater amount of heat will first be released (the steam condenses), one hundred degree water is formed, after which the same amount will be added as this water cools.
The phenomenon of a substance changing from a liquid to a gaseous state is called vaporization. Vaporization can be carried out in the form of two processes: evaporation And
Evaporation
Evaporation occurs from the surface of a liquid at any temperature. Thus, puddles dry out at 10 °C, 20 °C, and 30 °C. Thus, evaporation is the process of transformation of a substance from a liquid state into a gaseous state, occurring from the surface of a liquid at any temperature.
From the point of view of the structure of matter, the evaporation of a liquid is explained as follows. Liquid molecules, participating in continuous movement, have different speeds. The fastest molecules, located at the boundary of the surface of water and air and having relatively high energy, overcome the attraction of neighboring molecules and leave the liquid. Thus, above the liquid is formed steam.
Since molecules that have greater internal energy fly out of a liquid during evaporation compared to the energy of the molecules remaining in the liquid, the average speed and average kinetic energy of the liquid molecules decrease and, consequently, the temperature of the liquid decreases.
Evaporation rate liquid depends on the type of liquid. Thus, the rate of evaporation of ether is greater than the rate of evaporation of water and vegetable oil. In addition, the rate of evaporation depends on the movement of air above the surface of the liquid. The proof can be that laundry dries faster in the wind than in a windless place under the same external conditions.
Evaporation rate depends on the temperature of the liquid. For example, water at 30°C evaporates faster than water at 10°C.
It is well known that water poured into a saucer will evaporate faster than water of the same mass poured into a glass. Therefore, it depends on the surface area of the liquid.
Condensation
The process of changing a substance from a gaseous state to a liquid state is called condensation.
The condensation process occurs simultaneously with the evaporation process. Molecules emitted from the liquid and located above its surface participate in chaotic motion. They collide with other molecules, and at some point in time their speeds can be directed towards the surface of the liquid, and the molecules will return to it.
If the vessel is open, then the evaporation process occurs faster than condensation, and the mass of liquid in the vessel decreases. The vapor formed above a liquid is called unsaturated .
If the liquid is in a closed vessel, then at first the number of molecules leaving the liquid will be greater than the number of molecules returning to it, but over time the vapor density above the liquid will increase so much that the number of molecules leaving the liquid will become equal to the number of molecules returning to it. In this case, a dynamic equilibrium of the liquid with its vapor occurs.
Vapor that is in a state of dynamic equilibrium with its liquid is called saturated steam .
If a vessel with a liquid containing saturated steam is heated, then initially the number of molecules leaving the liquid will increase and will be greater than the number of molecules returning to it. Over time, equilibrium will be restored, but the density of the vapor above the liquid and, accordingly, its pressure will increase.
All substances have three states of aggregation - solid, liquid and gaseous, which appear under special conditions.
Definition 1
Phase transition is the transition of a substance from one state to another.
Examples of such a process are condensation and evaporation.
If you create certain conditions, you can turn any real gas (for example, nitrogen, hydrogen, oxygen) into a liquid. To do this, it is necessary to lower the temperature below a certain minimum, called the critical temperature. It is designated T to r. So, for nitrogen the value of this parameter is 126 K, for water – 647.3 K, for oxygen – 154.3 K. When maintaining room temperature, water can maintain both a gaseous and liquid state, while nitrogen and oxygen can only remain gaseous.
Definition 2
Evaporation- This is the phase transition of a substance into a gaseous state from a liquid.
The molecular kinetic theory explains this process by the gradual movement from the surface of the liquid of those molecules whose kinetic energy is greater than the energy of their connection with the rest of the molecules of the liquid substance. Due to evaporation, the average kinetic energy of the remaining molecules decreases, which, in turn, leads to a decrease in the temperature of the liquid if an additional source of external energy is not supplied to it.
Definition 3
Condensation is a phase transition of a substance from a gaseous state to a liquid state (the process reverse to evaporation).
During condensation, the vapor molecules return back to the liquid state.
Figure 3. 4 . 1 . Model of evaporation and condensation.
If a vessel containing a liquid or gas is clogged, then its contents may be in dynamic equilibrium, i.e. the speed of the condensation and evaporation processes will be the same (as many molecules will evaporate from the liquid as are returned back from the vapor). This system is called two-phase.
Definition 4
Saturated steam is a vapor that is in a state of dynamic equilibrium with its liquid.
There is a relationship between the number of molecules evaporating from the surface of a liquid in one second and the temperature of that liquid. The speed of the condensation process depends on the concentration of steam molecules and the speed of their thermal movement, which, in turn, is also directly dependent on temperature. Therefore, we can conclude that when a liquid and its vapor are in equilibrium, the concentration of molecules will be determined by the equilibrium temperature. As the temperature rises, a high concentration of vapor molecules is required so that evaporation and condensation become equal in speed.
Since, as we have already found out, concentration and temperature will determine the pressure of the vapor (gas), we can formulate the following statement:
Definition 5
The saturated vapor pressure p 0 of a certain substance does not depend on volume, but is directly dependent on temperature.
It is for this reason that isotherms of real gases on a plane include horizontal fragments that correspond to a two-phase system.
Figure 3. 4 . 2. Isotherms of real gas. Region I is liquid, region I I is a two-phase system “liquid + saturated vapor”, region I I I is a gaseous substance. K – critical point.
If the temperature rises, both the saturated vapor pressure and its density will increase, but the density of the liquid, on the contrary, will decrease due to thermal expansion. When the critical temperature for a given substance is reached, the densities of the liquid and gas are equalized; after passing this point, the physical differences between saturated vapor and liquid disappear.
Let's take saturated steam and compress it isothermally at T< T к р. Его давление будет постепенно возрастать, пока не сравняется с давлением насыщенного пара. Постепенно на дне сосуда появится жидкость, и между ней и ее насыщенным паром возникнет динамическое равновесие. По мере уменьшения объема будет происходить конденсация все большей части пара при неизменном давлении (на изотерме это состояние соответствует горизонтальному участку). После того, как весь пар перейдет в жидкое состояние, давление начнет резко увеличиваться при дальнейшем уменьшении объема, поскольку жидкость сжимается слабо.
It is not necessary to go through a two-phase region to make the transition from gas to liquid. The process can also be carried out bypassing the critical point. In the image, this option is shown using a broken line A B C.
Figure 3. 4 . 3. Isotherm model of real gas.
The air we breathe always contains water vapor at some pressure. This pressure is usually less than the saturated vapor pressure.
Definition 6
Relative humidity is the ratio of partial pressure to saturated water vapor pressure.
This can be written as a formula:
φ = p p 0 · 100 % .
To describe unsaturated steam, it is also permissible to use the equation of state of an ideal gas, taking into account the usual restrictions for real gas: not too high a vapor pressure (p ≤ (10 6 - 10 7) Pa) and a temperature higher than the value determined for each specific substance.
The ideal gas laws apply to describe saturated steam. However, the pressure for each temperature must be determined from the equilibrium curve for a given substance.
The higher the temperature, the higher the saturated vapor pressure. This dependence cannot be derived from the ideal gas laws. Assuming a constant concentration of molecules, the gas pressure will constantly increase in direct proportion to the temperature. If the vapor is saturated, then with increasing temperature not only the concentration will increase, but also the average kinetic energy of the molecules. It follows from this that the higher the temperature, the faster the saturated vapor pressure increases. This process occurs faster than the increase in pressure of an ideal gas, provided that the concentration of molecules in it remains constant.
What is boiling
We indicated above that evaporation occurs mainly from the surface, but it can also occur from the main volume of the liquid. Any liquid substance includes small gas bubbles. If the external pressure (i.e., the gas pressure in them) is equalized with the pressure of the saturated vapor, then the liquid inside the bubbles will evaporate, and they will begin to fill with steam, expand and float to the surface. This process is called boiling. Thus, the boiling point depends on the external pressure.
Definition 7
The liquid begins to boil at a temperature at which the external pressure and the pressure of its saturated vapors are equal.
If the atmospheric pressure is normal, then a temperature of 100 ° C is needed to boil water. At this temperature, the pressure of saturated water vapor will be equal to 1 a t m. If we boil water in the mountains, then due to a decrease in atmospheric pressure, the boiling point will drop to 70 ° C .
A liquid can only boil in an open vessel. If it is hermetically sealed, the balance between the liquid and its saturated vapor will be disrupted. You can find out the boiling point at different pressures using the equilibrium curve.
The image above shows the processes of phase transitions - condensation and evaporation using an isotherm of a real gas. This diagram is incomplete, since a substance can also take on a solid state. Achieving thermodynamic equilibrium between the phases of a substance at a given temperature is possible only at a certain pressure in the system.
Definition 8
Phase equilibrium curve is the relationship between equilibrium pressure and temperature.
An example of such a relationship could be the equilibrium curve between liquid and saturated vapor. If we construct curves that display the equilibrium between the phases of one substance on a plane, then we will see certain areas that correspond to different aggregate states of the substance - liquid, solid, gaseous. Curves plotted in a coordinate system are called phase diagrams.
Figure 3. 4 . 4 . Typical phase diagram of a substance. K – critical point, T – triple point. Region I is a solid, region I I is a liquid, region I I I is a gaseous substance.
The equilibrium between the gaseous and solid phases of a substance is reflected by the so-called sublimation curve (in the figure it is designated as 0 T), between vapor and liquid - by the evaporation curve, which ends at the critical point. The equilibrium curve between a liquid and a solid is called a melting curve.
Definition 9
Triple point– this is the point at which all equilibrium curves converge, i.e. All phases of matter are possible.
Many substances reach the triple point at a pressure less than 1 a t m ≈ 10 5 Pa. They melt when heated at atmospheric pressure. So, near water the triple point has coordinates T t r = 273.16 K, p t r = 6.02 10 2 P a. It is on this that the Kelvin absolute temperature scale is based.
For some substances, the triple point is reached at pressures above 1 atm.
Example 1
For example, carbon dioxide requires a pressure of 5.11 a t m and a temperature T tr = 216.5 K. If the pressure is equal to atmospheric, then to maintain it in a solid state, a low temperature is needed, and the transition to a liquid state becomes impossible. Carbon dioxide in equilibrium with its vapor at atmospheric pressure is called dry ice. This substance is not capable of melting, but can only evaporate (sublimate).
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