>>Physics: Dependence of saturation vapor pressure on temperature. Boiling

The liquid doesn't just evaporate. It boils at a certain temperature.
Saturated vapor pressure versus temperature. The state of saturated steam, as experience shows (we talked about this in the previous paragraph), is approximately described by the equation of state of an ideal gas (10.4), and its pressure is determined by the formula

As the temperature rises, the pressure rises. Because Saturated vapor pressure does not depend on volume, therefore, it depends only on temperature.
However, dependence r n.p. from T, found experimentally, is not directly proportional, as in an ideal gas at constant volume. With increasing temperature, the pressure of a real saturated vapor increases faster than the pressure of an ideal gas ( fig.11.1, section of the curve AB). This becomes obvious if we draw the isochores of an ideal gas through the points BUT and AT(dashed lines). Why is this happening?

When a liquid is heated in a closed vessel, part of the liquid turns into vapor. As a result, according to formula (11.1) saturated vapor pressure increases not only due to an increase in the temperature of the liquid, but also due to an increase in the concentration of molecules (density) of the vapor. Basically, the increase in pressure with increasing temperature is determined precisely by the increase in concentration. The main difference in the behavior of an ideal gas and saturated steam is that when the temperature of the vapor in a closed vessel changes (or when the volume changes at a constant temperature), the mass of the vapor changes. The liquid partially turns into vapor, or, conversely, the vapor partially condenses. Nothing like this happens with an ideal gas.
When all the liquid has evaporated, the vapor will cease to be saturated upon further heating, and its pressure at constant volume will increase in direct proportion to the absolute temperature (see Fig. fig.11.1, section of the curve Sun).
. As the temperature of the liquid increases, the rate of evaporation increases. Finally, the liquid begins to boil. When boiling, rapidly growing vapor bubbles form throughout the volume of the liquid, which float to the surface. The boiling point of a liquid remains constant. This is because all the energy supplied to the liquid is spent on turning it into steam. Under what conditions does boiling begin?
The liquid always contains dissolved gases that are released on the bottom and walls of the vessel, as well as on dust particles suspended in the liquid, which are the centers of vaporization. The liquid vapors inside the bubbles are saturated. As the temperature increases, the vapor pressure increases and the bubbles increase in size. Under the action of the buoyant force, they float up. If the upper layers of the liquid have a lower temperature, then vapor condenses in these layers in the bubbles. The pressure drops rapidly and the bubbles collapse. The collapse is so fast that the walls of the bubble, colliding, produce something like an explosion. Many of these microexplosions create a characteristic noise. When the liquid warms up enough, the bubbles stop collapsing and float to the surface. The liquid will boil. Watch the kettle on the stove carefully. You will find that it almost stops making noise before boiling.
The dependence of saturation vapor pressure on temperature explains why the boiling point of a liquid depends on the pressure on its surface. A vapor bubble can grow when the pressure of the saturated vapor inside it slightly exceeds the pressure in the liquid, which is the sum of the air pressure on the surface of the liquid (external pressure) and the hydrostatic pressure of the liquid column.
Let us pay attention to the fact that the evaporation of a liquid occurs at temperatures lower than the boiling point, and only from the surface of the liquid; during boiling, the formation of vapor occurs throughout the entire volume of the liquid.
Boiling begins at a temperature at which the saturation vapor pressure in the bubbles is equal to the pressure in the liquid.
The greater the external pressure, the higher the boiling point. So, in a steam boiler at a pressure reaching 1.6 10 6 Pa, water does not boil even at a temperature of 200°C. AT medical institutions in hermetically sealed vessels - autoclaves ( fig.11.2) boiling of water also occurs at high blood pressure. Therefore, the boiling point of the liquid is much higher than 100°C. Autoclaves are used for sterilization surgical instruments and etc.

And vice versa, reducing the external pressure, we thereby lower the boiling point. By pumping out air and water vapor from the flask, you can make the water boil at room temperature (fig.11.3). As you climb mountains, atmospheric pressure decreases, so the boiling point decreases. At an altitude of 7134 m (Lenin Peak in the Pamirs), the pressure is approximately 4 10 4 Pa ​​(300 mm Hg). Water boils there at about 70°C. It is impossible to cook meat in these conditions.

Each liquid has its own boiling point, which depends on the pressure of its saturated vapor. The higher the saturated vapor pressure, the lower the boiling point of the liquid, since at lower temperatures the saturated vapor pressure becomes equal to atmospheric pressure. For example, at a boiling point of 100 ° C, the pressure of saturated water vapor is 101,325 Pa (760 mm Hg), and mercury vapor is only 117 Pa (0.88 mm Hg). Mercury boils at 357°C at normal pressure.
A liquid boils when its saturated vapor pressure becomes equal to the pressure inside the liquid.

???
1. Why does the boiling point increase with increasing pressure?
2. Why is it essential for boiling to increase the pressure of saturated vapor in the bubbles, and not to increase the pressure of the air present in them?
3. How to make a liquid boil by cooling the vessel? (This is a tricky question.)

G.Ya.Myakishev, B.B.Bukhovtsev, N.N.Sotsky, Physics Grade 10

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To regulate the evaporation pressure, use the KVP regulator, which is installed on the suction line downstream of the evaporator (Fig. 6.13).

In addition to its main function, the evaporating pressure regulator provides protection in the event of a strong drop in evaporating pressure in order to prevent freezing of the cooled water in the heat exchange path of the evaporators of water chillers.

The regulator works as follows: when the pressure rises above the set pressure, the regulator opens, and when the pressure drops below the set value, it closes. The control signal is only the pressure at the inlet to the regulator.

In installations with several evaporators and operating at different evaporation pressures, the regulator is installed after the evaporator, in which the pressure is highest. To avoid condensation of the refrigerant during stops, a non-return valve is mounted on the suction line immediately after the evaporator with minimum pressure. In installations with parallel evaporators and a common compressor, a regulator is installed on the suction line to maintain the same pressure in the evaporators.

In addition to this type of regulator, the evaporation pressure is stabilized using electronic control systems for one or more refrigerating chambers, cabinets, etc., providing high accuracy in maintaining the set temperature (± 0.5 K) in a wide range of cooling capacity - from 10 to 100% of the nominal value .

8. Performance regulators.

The capacity controllers (fig. 6.14) help the compressor's cooling capacity to adapt to the change in the heat load on the evaporator in installations with a very low refrigerant charge. They avoid low suction pressure and useless starts.

As the heat load on the evaporator decreases, the suction pressure drops, causing a vacuum in the circuit, which leads to the danger of moisture entering the installation. When the suction pressure falls below the set value, the regulator opens, causing a certain volume of hot gases to pass from the discharge to the suction. As a result, the suction pressure increases and the cooling capacity decreases. The regulator reacts only to the pressure in the suction line, i.e. at the exit from it.

9. Starting regulators.

The start regulators make it possible to avoid running and starting the compressor at too high suction pressure values, which occurs after a long stop of the machine or after defrosting the evaporator.

The starting regulator KVL belongs to the type of throttling pressure regulators "after itself". It maintains a constant pressure in the suction line between the regulator and the compressor and unloads the compressor at start-up.

Regulator inlet pressure acts on the bellows from below and on the valve plate from above. Since the effective area of ​​the bellows is equivalent to the area of ​​the orifice, the inlet pressure is neutralized. The outlet pressure (in the crankcase) acts on the valve plate from below, counteracting the tension force of the adjustable spring. These two forces are the acting forces of the regulator. When the regulated pressure at the outlet (in the crankcase) decreases, the valve opens, passing refrigerant vapor into the compressor. For refrigeration units high capacity, parallel mounting of KVL start-up controllers is possible. In this case, the regulators are selected from the condition of the same pressure drop in each pipeline and equivalent performance.

The regulator is adjusted to the maximum values, without exceeding, however, the values ​​recommended by the manufacturer for the compressor or condensing unit. The setting is carried out according to the readings of the pressure gauge on the suction line of the compressor.

The start regulator is installed on the suction line between the evaporator and the compressor (Fig. 6.15).

In this regulator, it is possible to connect a vapor extraction line through a manometric outlet on the inlet pipe having a 1/4" bore diameter. With this method of regulation, vapor extraction is carried out "after itself".

The choice of a starting regulator is determined by five main indicators:

The type of refrigerant

system performance,

suction design pressure,

Maximum design pressure,

pressure drop in the regulator.

The difference between design and maximum design suction pressure determines how long the valve will open. The pressure drop across the regulator is an important factor, as pressure losses in the suction line affect machine performance. Therefore, pressure drop across the regulator must be kept to a minimum. Typically, in low temperature refrigeration systems, the pressure drop is 3...7 kPa. The maximum pressure drop for most refrigeration systems is 14 kPa.

At the maximum opening of the valve, the regulator, on the one hand, provides maximum performance, and on the other hand, causes large pressure losses, which reduces the performance of the system. Therefore, the pressure drop across the regulator must be kept as low as possible.

The process of intense evaporation of a liquid begins at a temperature when the vapor pressure of the liquid exceeds the external pressure of the gas atmosphere above the liquid. At the boiling point, the formation of steam occurs in the entire mass of the liquid and flows at an almost constant temperature until the complete transition of the liquid (single-component) and vapor. By artificially lowering the pressure, it is possible to make the liquid boil at lower temperatures, which is widely used in technology, since it is easier to find suitable material for hardware. Modern vacuum technology has at its disposal powerful rotary pumps capable of creating a vacuum at which the residual pressure does not exceed 0.001 mmHg, and jet diffusion pumps that create a vacuum of up to 10v-7-10v-8 mmHg. Art.
Vacuum distillation is used to obtain high purity metals; Zn, Cd, Mg, Ca, etc. Usually they work at pressures slightly higher than the vapor pressure of the distilled metal at its melting point. Then, by distilling the liquid metal, a solid condensate is obtained, which makes it possible to apply very simple design distillation apparatus shown in Fig. 24. The device is a cylinder, in the lower part of which there is a vessel with liquid distilled metal. Vapors are condensed in the upper part of the cylinder on a special composite metal cylinder(capacitor) in the form of a crystalline crust, which, after the end of the process, is removed together with the condenser. Before heating the metal, first, the air is pumped out of the device with a vacuum pump, and then from time to time the vacuum is restored, which changes due to air leakage from the outside through the leakage of the equipment. If the apparatus is sufficiently hermetic, then during the distillation process, since non-condensable gases are not released, the constant operation of the vacuum pump is not necessary.

The described device is extremely simple, it is made of steel or heat-resistant metal alloys. What is especially important, its cover and all sealing - sealing parts are cooled with water, i.e. they operate at room temperature, which allows the use of very advanced sealants - rubber, vacuum putties, etc. The use of vacuum allows cleaning by distillation at relatively low temperatures (700 -900 °) such chemically active and very aggressive metals as calcium, magnesium, barium, the distillation of which at atmospheric pressure is not feasible due to the impossibility of selecting material for the equipment.
Let us consider the features of the evaporation process in vacuum.
The state diagram liquid - vapor with decreasing pressure has the same character as the diagrams for atmospheric pressure, only the lines of liquid and vapor move to the region of lower temperatures. It follows that the efficiency of separation of the components during the evaporation of their solution in vacuum is approximately the same as at atmospheric pressure, but is carried out at lower temperatures; the temperature is lower the deeper the applied vacuum. A feature of vacuum operation is the absence of entrainment of small liquid droplets along with vapors, which is always observed when operating under atmospheric pressure. During the rapid boiling of the liquid, the bursting bubbles of the steam rising from the depth of the liquid give splashes, which are carried away by the vapors into the condenser and pollute the distillate. In a vacuum (deep enough), spattering does not occur, since the boiling process is fundamentally different from boiling at atmospheric pressure. In a vacuum, the formation of steam occurs only on the surface of the liquid, bubbles do not form inside the liquid, the surface is calm, does not boil, therefore, splashes cannot occur. Therefore, vacuum distillation produces a purer distillate than atmospheric distillation.
Let us use an example to show the peculiarity of the process of boiling in vacuum. Let in one case water in a vessel with a layer depth of 250 mm boil at atmospheric pressure (760 mm Hg). Then the steam released from the surface of the water, in order to overcome the external pressure, must have atmospheric pressure (760 mm Hg), which develops at a water surface temperature of 100 °. The vapor bubble formed at the bottom of the vessel must have a higher pressure, since, in addition to the pressure of the atmosphere, it needs to overcome the hydrostatic pressure of a water column 250 mm high, which corresponds to an excess pressure of 18 mm Hg. Art. Thus, the steam released from the bottom of the vessel must have a pressure of 760 + 18 = 778 mm Hg. Art., which corresponds to the temperature of the water at the bottom of the vessel 100.6 °. Such a small overheating of water at the bottom (0.6°) is quite real, and the boiling process proceeds in such a way that steam is formed in the entire mass of the layer. The water boils vigorously and splashes as the bubbles break on the surface.
Now consider boiling the same layer of water in a vacuum of 4.58 mm Hg. Art. For boiling, the surface layer of water must have a temperature of 0 °, at which the saturated vapor pressure is 4.58 mm Hg. Art. The bubble that forms at the bottom must overcome the hydrostatic pressure of a water column of 250 mm, which corresponds to a pressure of 18 mm Hg. Art., and have a total pressure of 4.58 + 18 = 22.58 mm Hg. Art. Water will have such a saturated vapor pressure at a temperature of ~ 23 °, i.e., in order for a vapor bubble to form at the bottom of the vessel, it is necessary to have a temperature of 23 ° at the bottom. It is impossible to obtain such a difference between the temperatures near the bottom and on the surface, since this will be prevented by convection currents. Consequently, bubbles will not form in the depth of the liquid layer and vaporization will be carried out only from the surface of the liquid.
Metal melts have high thermal conductivity, which prevents local overheating of the liquid and, consequently, boiling with the formation of bubbles.
Until the pressure in the device becomes very small, an exchange of molecules takes place between the surface of the liquid and the vapor and a mobile liquid-vapor equilibrium is established. A conventional gas stream of vapor flows to the condenser and the results of the distillation process are determined by the liquid-vapor state diagram.
If the pressure in the instrument is so low that the mean free path of the molecules becomes more sizes device, the nature of the distillation process changes radically.
Under these conditions, there is no exchange of molecules between vapors and liquid, a mobile liquid-vapor equilibrium is not established, and the liquid-vapor state diagram does not describe the evaporation process. Ordinary gas eschar between evaporator and condenser. He is formed, the vapor molecules separated from the surface of the liquid follow a straight path, without colliding with other molecules, they fall on the cold surface of the condenser and remain there - they condense; the evaporation process is completely irreversible and has the character of molecular evaporation. The result of distillation is determined by the evaporation rate, which depends on the type of the evaporated substance and temperature and is independent of the external pressure in the system, if this pressure is sufficiently low. The evaporation rate under these conditions can be calculated using the Langmuir formula:

Taking as the rate of evaporation the mass of a substance evaporating per second from a unit surface, expressing the vapor pressure p in millimeters mercury column and replacing the quantities R and π with their numerical values, we obtain equation (III, 13) in a different form, convenient for practical calculations:

In molecular evaporation, substances with the same vapor pressure can be separated if their molecular weights are different, as proven by experiments on isotope separation.

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It is clear from the above reasoning that the boiling point of a liquid must depend on the external pressure. Observations confirm this.

The greater the external pressure, the higher the boiling point. So, in a steam boiler at a pressure reaching 1.6 10 6 Pa, water does not boil even at a temperature of 200 °C. In medical institutions, boiling water in hermetically sealed vessels - autoclaves (Fig. 6.11) also occurs at elevated pressure. Therefore, the boiling point is much higher than 100 °C. Autoclaves are used to sterilize surgical instruments, dressings, etc.

Conversely, by reducing the external pressure, we thereby lower the boiling point. Under the bell of the air pump, you can make water boil at room temperature (Fig. 6.12). As you climb mountains, atmospheric pressure decreases, so the boiling point decreases. At an altitude of 7134 m (Lenin Peak in the Pamirs), the pressure is approximately 4 10 4 Pa ​​(300 mm Hg). Water boils there at about 70°C. It is impossible to cook, for example, meat in these conditions.

Figure 6.13 shows the dependence of the boiling point of water on external pressure. It is easy to see that this curve is also a curve expressing the dependence of saturated water vapor pressure on temperature.

The difference in boiling points of liquids

Each liquid has its own boiling point. The difference in the boiling points of liquids is determined by the difference in the pressure of their saturated vapors at the same temperature. For example, ether vapor already at room temperature has a pressure greater than half atmospheric pressure. Therefore, in order for the ether vapor pressure to become equal to atmospheric, a slight increase in temperature (up to 35 ° C) is needed. In mercury, saturated vapors have a very negligible pressure at room temperature. The vapor pressure of mercury becomes equal to atmospheric only with a significant increase in temperature (up to 357 ° C). It is at this temperature, if the external pressure is 105 Pa, that mercury boils.

The difference in the boiling points of substances is of great use in technology, for example, in the separation of petroleum products. When oil is heated, its most valuable, volatile parts (gasoline) evaporate first of all, which can thus be separated from “heavy” residues (oils, fuel oil).

A liquid boils when its saturated vapor pressure equals the pressure inside the liquid.

§ 6.6. Heat of vaporization

Is energy required to turn liquid into vapor? Probably yes! Is not it?

We noted (see § 6.1) that the evaporation of a liquid is accompanied by its cooling. To maintain the temperature of the evaporating liquid unchanged, heat must be supplied to it from the outside. Of course, heat itself can be transferred to liquid from surrounding bodies. So, the water in the glass evaporates, but the temperature of the water, which is somewhat lower than the temperature of the surrounding air, remains unchanged. Heat is transferred from the air to the water until all the water has evaporated.

To keep water (or any other liquid) boiling, heat must also be continuously supplied to it, for example, by heating it with a burner. In this case, the temperature of the water and the vessel does not rise, but a certain amount of steam is formed every second.

Thus, in order to convert a liquid into vapor by evaporation or by boiling, an influx of heat is required. The amount of heat required to convert a given mass of liquid into vapor at the same temperature is called the heat of vaporization of that liquid.

What is the energy supplied to the body used for? First of all, to increase it internal energy during the transition from a liquid to a gaseous state: in this case, the volume of a substance increases from the volume of liquid to the volume of saturated vapor. Consequently, the average distance between molecules increases, and hence their potential energy.

In addition, when the volume of a substance increases, work is done against the forces of external pressure. This part of the heat of vaporization at room temperature is usually a few percent of the total heat of vaporization.

The heat of vaporization depends on the type of liquid, its mass and temperature. The dependence of the heat of vaporization on the type of liquid is characterized by a value called the specific heat of vaporization.

The specific heat of vaporization of a given liquid is the ratio of the heat of vaporization of a liquid to its mass:

(6.6.1)

where r- specific heat of vaporization of the liquid; t- mass of liquid; Q n is its heat of vaporization. The SI unit for specific heat of vaporization is the joule per kilogram (J/kg).

The specific heat of vaporization of water is very high: 2.256 10 6 J/kg at a temperature of 100 °C. For other liquids (alcohol, ether, mercury, kerosene, etc.), the specific heat of vaporization is 3-10 times less.

States of matter

Iron vapor and solid air

Isn't it a strange combination of words? However, this is not nonsense at all: both iron vapor and solid air exist in nature, but not under ordinary conditions.

What conditions are we talking about? The state of matter is determined by two circumstances: temperature and pressure.

Our life takes place in relatively little changing conditions. Air pressure fluctuates within a few percent around one atmosphere; air temperature, say, in the Moscow area lies in the range from -30 to + 30 ° C; in the absolute temperature scale, in which the lowest possible temperature (-273 ° C) is taken as zero; this interval will look less impressive: 240-300 K, which is also only ±10% of the average value.

It is quite natural that we are accustomed to these ordinary conditions, and therefore, when we say simple truths like: "iron is a solid, air is a gas," etc., we forget to add: "under normal conditions."

If iron is heated, it first melts and then evaporates. If the air is cooled, it will first turn into a liquid, and then solidify.

Even if the reader has never met with iron vapor and solid air, he will probably easily believe that any substance, by changing the temperature, can be obtained in solid, liquid, and gaseous states, or, as they say, in solid , liquid or gaseous phases.

It is easy to believe in this because one substance, without which life on Earth would be impossible, everyone observed both in the form of a gas, and as a liquid, and in the form of a solid body. We are, of course, talking about water.

What are the conditions under which a substance changes from one state to another?

Boiling

If we lower the thermometer into the water that is poured into the kettle, turn on the electric stove and monitor the mercury of the thermometer, we will see the following: almost immediately the level of mercury will creep up. It's already 90, 95, finally 100°C. The water boils, and at the same time the rise of mercury stops. The water has been boiling for many minutes, but the level of mercury does not change. Until all the water boils away, the temperature will not change (Fig. 4.1).

Rice. 4.1

Where does the heat go if the temperature of the water does not change? The answer is obvious. The process of turning water into steam requires energy.

Let's compare the energy of a gram of water and a gram of steam formed from it. Vapor molecules are farther apart than water molecules. It is clear that because of this, the potential energy of water will differ from potential energy pair.

The potential energy of attracted particles decreases as they approach each other. Therefore, the energy of steam is greater than the energy of water, and the transformation of water into steam requires energy. This excess of energy is communicated by an electric stove to boiling water in a kettle.

The energy needed to turn water into steam; called the heat of vaporization. It takes 539 calories to turn 1 g of water into steam (this is the figure for a temperature of 100°C).

If 539 cal goes to 1 g, then 18 * 539 \u003d 9700 cal will be spent on 1 mole of water. This amount of heat must be expended to break intermolecular bonds.

You can compare this figure with the amount of work required to break intramolecular bonds. In order to split 1 mole of water vapor into atoms, about 220,000 calories are required, that is, 25 times more energy. This directly proves the weakness of the forces that bind molecules to each other, compared with the forces that pull atoms together into a molecule.

Boiling temperature versus pressure

The boiling point of water is 100°C; one might think that this is an inherent property of water, that water, wherever and under what conditions it is, will always boil at 100 ° C.

But this is not so, and the inhabitants of high-mountain villages are well aware of this.

Near the top of Elbrus there is a house for tourists and a scientific station. Beginners sometimes wonder "how difficult it is to boil an egg in boiling water" or "why boiling water does not burn." Under these conditions, they are told that water boils at the top of Elbrus already at 82°C.

What is the matter here? What physical factor interferes with the phenomenon of boiling? What is the significance of altitude?

This physical factor is the pressure acting on the surface of the liquid. You do not need to climb to the top of the mountain to check the validity of what has been said.

By placing heated water under the bell and pumping air in or out of it, one can be convinced that the boiling point rises with increasing pressure and falls with decreasing pressure.

Water boils at 100°C only at a certain pressure - 760 mm Hg. Art. (or 1 atm).

The boiling point versus pressure curve is shown in fig. 4.2. At the top of Elbrus, the pressure is 0.5 atm, and this pressure corresponds to a boiling point of 82 ° C.

Rice. 4.2

But water boiling at 10-15 mm Hg. Art., you can freshen up in hot weather. At this pressure, the boiling point will drop to 10-15°C.

You can even get "boiling water", which has the temperature of freezing water. To do this, you will have to reduce the pressure to 4.6 mm Hg. Art.

An interesting picture can be observed if you place an open vessel with water under the bell and pump out the air. Pumping will make the water boil, but boiling requires heat. There is nowhere to take it from, and the water will have to give up its energy. The temperature of the boiling water will begin to drop, but as the pumping continues, so will the pressure. Therefore, the boiling will not stop, the water will continue to cool and eventually freeze.

Such a boil cold water occurs not only when pumping air. For example, when a ship's propeller rotates, the pressure in a rapidly moving metal surface the water layer drops strongly and the water in this layer boils, i.e., numerous bubbles filled with steam appear in it. This phenomenon is called cavitation (from the Latin word cavitas - cavity).

By lowering the pressure, we lower the boiling point. What about increasing it? A graph like ours answers this question. A pressure of 15 atm can delay the boiling of water, it will only start at 200°C, and a pressure of 80 atm will make water boil only at 300°C.

So, a certain external pressure corresponds to a certain boiling point. But this statement can also be "turned over", saying this: each boiling point of water corresponds to its own specific pressure. This pressure is called vapor pressure.

The curve depicting the boiling point as a function of pressure is also the curve of vapor pressure as a function of temperature.

Figures plotted on a boiling point graph (or vapor pressure graph) show that vapor pressure changes very rapidly with temperature. At 0°C (i.e., 273 K), the vapor pressure is 4.6 mm Hg. Art., at 100 ° C (373 K) it is equal to 760 mm Hg. Art., i.e. increases by 165 times. When the temperature doubles (from 0 ° C, i.e. 273 K, to 273 ° C, i.e. 546 K), the vapor pressure increases from 4.6 mm Hg. Art. up to almost 60 atm, i.e., about 10,000 times.

Therefore, on the contrary, the boiling point changes rather slowly with pressure. When the pressure is doubled from 0.5 atm to 1 atm, the boiling point increases from 82°C (355 K) to 100°C (373 K) and when the pressure is doubled from 1 to 2 atm, from 100°C (373 K) to 120°C (393 K).

The same curve that we are now considering also controls the condensation (thickening) of steam into water.

Steam can be converted to water by either compression or cooling.

Both during boiling and during condensation, the point will not move off the curve until the conversion of steam to water or water to steam is complete. This can also be formulated as follows: under the conditions of our curve, and only under these conditions, the coexistence of liquid and vapor is possible. If at the same time no heat is added or taken away, then the quantities of vapor and liquid in a closed vessel will remain unchanged. Such vapor and liquid are said to be in equilibrium, and a vapor in equilibrium with its liquid is said to be saturated.

The curve of boiling and condensation, as we see, has another meaning: it is the equilibrium curve of liquid and vapor. The equilibrium curve divides the diagram field into two parts. To the left and upwards (toward higher temperatures and lower pressures) is the region of the steady state of steam. To the right and down - the region of the stable state of the liquid.

The vapor-liquid equilibrium curve, i.e., the dependence of the boiling point on pressure or, what is the same, vapor pressure on temperature, is approximately the same for all liquids. In some cases, the change may be somewhat more abrupt, in others - somewhat slower, but always the vapor pressure increases rapidly with increasing temperature.

We have used the words "gas" and "steam" many times. These two words are pretty much the same. We can say: water gas is the vapor of water, gas oxygen is the vapor of an oxygen liquid. Nevertheless, some habit has developed in the use of these two words. Since we are accustomed to a certain relatively small temperature range, we usually apply the word "gas" to those substances whose vapor pressure at ordinary temperatures is above atmospheric pressure. On the contrary, we speak of a vapor when, at room temperature and atmospheric pressure, the substance is more stable in the form of a liquid.

Evaporation

Boiling is a fast process, and in a short time there is no trace of boiling water, it turns into steam.

But there is another phenomenon of the transformation of water or other liquid into steam - this is evaporation. Evaporation occurs at any temperature, regardless of pressure, which under normal conditions is always close to 760 mm Hg. Art. Evaporation, unlike boiling, is a very slow process. The bottle of cologne we forgot to close will be empty in a few days; more time o a saucer with water will stand, but sooner or later it will turn out to be dry.

Air plays an important role in the evaporation process. By itself, it does not prevent water from evaporating. As soon as we open the surface of the liquid, water molecules will begin to move into the nearest layer of air.

The vapor density in this layer will increase rapidly; after a short time, the vapor pressure will become equal to the elasticity characteristic of the temperature of the medium. In this case, the vapor pressure will be exactly the same as in the absence of air.

The transition of vapor into air does not, of course, mean an increase in pressure. The total pressure in the space above the water surface does not increase, only the share in this pressure that is taken on by steam increases, and, accordingly, the proportion of air that is displaced by steam decreases.

Above the water there is steam mixed with air, above there are layers of air without steam. They will inevitably mix. Water vapor will continuously move to higher layers, and in its place, air will flow into the lower layer, which does not contain water molecules. Therefore, in the layer closest to the water, places will always be freed up for new water molecules. The water will continuously evaporate, maintaining the water vapor pressure at the surface equal to the elasticity, and the process will continue until the water has completely evaporated.

We started with the cologne and water example. It is well known that they evaporate at different rates. Ether evaporates exceptionally quickly, alcohol rather quickly, and water much more slowly. We will immediately understand what is the matter if we find in the reference book the values ​​of the vapor pressure of these liquids, say, at room temperature. Here are the numbers: ether - 437 mm Hg. Art., alcohol - 44.5 mm Hg. Art. and water - 17.5 mm Hg. Art.

The greater the elasticity, the more vapor in the adjacent layer of air and the faster the liquid evaporates. We know that vapor pressure increases with temperature. It is clear why the rate of evaporation increases with heating.

The rate of evaporation can also be influenced in another way. If we want to help evaporation, we must quickly remove the vapor from the liquid, i.e., speed up the mixing of air. That is why evaporation is greatly accelerated by blowing the liquid. Water, although it has a relatively small vapor pressure, will disappear rather quickly if the saucer is placed in the wind.

It is understandable, therefore, why a swimmer who comes out of the water feels cold in the wind. The wind accelerates the mixing of air with steam and, therefore, accelerates evaporation, and the heat for evaporation is forced to give up the human body.

A person's well-being depends on whether there is a lot or a little water vapor in the air. Both dry and humid air are unpleasant. Humidity is considered normal when it is 60%. This means that the density of water vapor is 60% of the density of saturated water vapor at the same temperature.

If moist air is cooled, then eventually the pressure of water vapor in it will be equal to the vapor pressure at this temperature. The steam will become saturated and, as the temperature drops further, it will begin to condense into water. Morning dew, moisturizing grass and leaves, appears just because of this phenomenon.

At 20°C, the density of saturated water vapor is about 0.00002 g/cm 3 . We will feel good if the air contains 60% of this number of water vapor - which means only a little more than one hundred thousandth of a gram in 1 cm 3.

Although this figure is small, it will lead to impressive amounts of steam for a room. It is easy to calculate that in a medium-sized room with an area of ​​12 m 2 and a height of 3 m, about a kilogram of water can "fit" in the form of saturated steam.

So, if you tightly close such a room and put an open barrel of water, then a liter of water will evaporate, no matter what the capacity of the barrel.

It is interesting to compare this result for water with the corresponding figures for mercury. At the same temperature of 20°C, the density of saturated mercury vapor is 10 -8 g/cm 3 .

In the room we have just discussed, no more than 1 g of mercury vapor will fit.

By the way, mercury vapor is very toxic, and 1 g of mercury vapor can seriously damage the health of any person. When working with mercury, care must be taken that even the smallest drop of mercury does not spill.

Critical temperature

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

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

However, some gases - hydrogen, nitrogen, oxygen - did not lend themselves to liquefaction. No matter how much the pressure was increased, they did not turn into a liquid. One might have thought that oxygen and other gases could not be liquid. They were classified as true, or permanent, gases.

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

Consider a liquid and a vapor in equilibrium and consider what happens to them as the boiling point rises and, of course, as the pressure rises accordingly. In other words, imagine that a point on the boiling graph moves up along the curve. It is clear that the liquid expands with increasing temperature and its density decreases. As for steam, an increase in the boiling point? of course, it contributes to its expansion, but, as we have already said, the saturation vapor pressure rises much faster than the boiling point. Therefore, the vapor density does not fall, but, on the contrary, increases rapidly with increasing boiling point.

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

Rice. 4.3

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

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

Rice. 4.4

The situation is completely different when the gas is compressed, the temperature of which is higher than the critical one. The compression process can again be depicted as an arrow going from bottom to top. But now this arrow does not cross the boiling curve. This means that during compression, the vapor will not condense, but will only continuously condense.

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

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

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

Thus, it is always possible, by varying the pressure and temperature, bypassing the critical point, to obtain vapor by continuous transition from liquid or liquid from vapor. Such a continuous transition does not require boiling or condensation.

Early attempts to liquefy gases such as oxygen, nitrogen, hydrogen were therefore unsuccessful because the existence of a critical temperature was not known. These gases have very low critical temperatures: nitrogen has -147°C, oxygen has -119°C, hydrogen has -240°C, or 33 K. The record holder is helium, its critical temperature is 4.3 K. Turn these gases into liquid can only be done in one way - it is necessary to reduce their temperature below the specified one.

Getting low temperatures

A significant decrease in temperature can be achieved in various ways. But the idea of ​​all methods is the same: we must force the body that we want to cool down to expend its internal energy.

How to do it? One way is to make the liquid boil without supplying heat from outside. To do this, as we know, it is necessary to reduce the pressure - to reduce it to the value of vapor pressure. The heat expended for boiling will be borrowed from the liquid and the temperature of the liquid and vapor, and with it the vapor pressure will fall. Therefore, in order for the boiling not to stop and to happen faster, air must be continuously pumped out of the vessel with the liquid.

However, there is a limit to the temperature drop during this process: the vapor pressure eventually becomes completely insignificant, and even the strongest pumping pumps cannot create the required pressure.

In order to continue lowering the temperature, it is possible, by cooling the gas with the resulting liquid, to turn it into a liquid with a lower boiling point.

Now the pumping process can be repeated with the second substance and thus lower temperatures can be obtained. If necessary, such a "cascade" method for obtaining low temperatures can be extended.

This is exactly what they did at the end of the last century; the liquefaction of gases was carried out in stages: ethylene, oxygen, nitrogen, hydrogen, substances with boiling points of -103, -183, -196 and -253°C, were successively converted into liquid. Having liquid hydrogen, you can also get the lowest boiling liquid - helium (-269 ° C). The neighbor on the "left" helped to get the neighbor on the "right".

The cascade cooling method is almost a hundred years old. In 1877 liquid air was obtained by this method.

In 1884-1885. liquid hydrogen was produced for the first time. Finally, after another twenty years, the last fortress was taken: in 1908, Kamerling-Onnes in the city of Leiden in Holland turned helium into a liquid - a substance with the lowest critical temperature. The 70th anniversary of this important scientific achievement was recently celebrated.

For many years the Leiden Laboratory was the only "low-temperature" laboratory. Now in all countries there are dozens of such laboratories, not to mention plants that produce liquid air, nitrogen, oxygen and helium for technical purposes.

The cascade method for obtaining low temperatures is now rarely used. AT technical installations to lower the temperature, another method is used to lower the internal energy of the gas: they force the gas to expand rapidly and do work at the expense of internal energy.

If, for example, air compressed to several atmospheres is put into an expander, then when the work of moving the piston or rotating the turbine is performed, the air will cool so sharply that it will turn into a liquid. Carbon dioxide, if it is quickly released from the cylinder, cools so sharply that it turns into "ice" on the fly.

Liquid gases are widely used in engineering. Liquid oxygen is used in explosive technology as a component of the fuel mixture in jet engines.

Air liquefaction is used in engineering to separate the gases that make up air.

In various fields of technology, it is required to work at liquid air temperature. But for many physical studies, this temperature is not low enough. Indeed, if we translate degrees Celsius into an absolute scale, we will see that the temperature of liquid air is about 1/3 of room temperature. Much more interesting for physics are "hydrogen" temperatures, ie, temperatures of the order of 14-20 K, and especially "helium" temperatures. The lowest temperature obtained when liquid helium is pumped out is 0.7 K.

Physicists have managed to come much closer to absolute zero. At present, temperatures exceeding absolute zero by only a few thousandths of a degree have been obtained. However, these ultra-low temperatures are obtained in ways that are not similar to those that we have described above.

AT last years low-temperature physics gave rise to a special branch of industry engaged in the production of equipment that makes it possible to maintain large volumes at a temperature close to absolute zero; power cables have been developed whose busbars operate at a temperature of less than 10 K.

Supercooled vapor and superheated liquid

At the transition of the boiling point, the vapor must condense, turn into a liquid. However,; It turns out that if the vapor does not come into contact with the liquid, and if the vapor is very pure, then it is possible to obtain a supercooled or supersaturated vapor - a vapor that should have become a liquid long ago.

Supersaturated steam is very unstable. Sometimes a push or a grain of steam thrown into space is enough to start a belated condensation.

Experience shows that the condensation of vapor molecules is greatly facilitated by the introduction of small foreign particles into the vapor. In dusty air, supersaturation of water vapor does not occur. Can cause condensation with puffs of smoke. After all, smoke is made up of small solid particles. Getting into the steam, these particles collect molecules around themselves and become centers of condensation.

So, although unstable, steam can exist in the temperature range adapted for the "life" of the liquid.

Can a liquid `live' in the region of vapor under the same conditions? In other words, is it possible to superheat a liquid?

It turns out you can. To do this, it is necessary to ensure that the molecules of the liquid do not break away from its surface. The radical remedy is to eliminate the free surface, that is, to place the liquid in a vessel where it would be compressed on all sides by solid walls. In this way, it is possible to achieve overheating of the order of several degrees, i.e., to move the point depicting the state of liquids to the right of the boiling curve (Fig. 4.4).

Overheating is a shift of a liquid into a vapor region, so overheating of a liquid can be achieved both by supplying heat and by reducing pressure.

The last way you can achieve amazing results. Water or other liquid, carefully freed from dissolved gases (this is not easy to do), is placed in a vessel with a piston that reaches the surface of the liquid. The vessel and piston must be wetted by liquid. If you now pull the piston towards you, then the water adhered to the bottom of the piston will follow it. But the layer of water, clinging to the piston, will pull the next layer of water, this layer will pull the underlying one, as a result, the liquid will stretch.

In the end, the column of water will break (it is the column of water, and not the water, that will come off the piston), but this will happen when the force per unit area reaches tens of kilograms. In other words, a negative pressure of tens of atmospheres is created in the liquid.

Even at low positive pressures, the vapor state of matter is stable. A liquid can be brought to a negative pressure. You can't imagine a more striking example of "overheating".

Melting

There is no such solid body that would resist an increase in temperature as much as necessary. Sooner or later a solid piece turns into a liquid; right, in some cases we will not be able to get to the melting point - chemical decomposition can occur.

As the temperature rises, the molecules move faster and faster. Finally, there comes a moment when maintaining order "among strongly" swung "molecules becomes impossible. The solid body melts. Tungsten has the highest melting point: 3380 ° C. Gold melts at 1063 ° C, iron at 1539 ° C. However, there are also fusible metals. Mercury, as is well known, melts already at a temperature of -39 ° C. Organic substances do not have high melting points. Naphthalene melts at 80 ° C, toluene - at -94.5 ° C.

It is not at all difficult to measure the melting point of a body, especially if it melts in the temperature range that is measured with an ordinary thermometer. It is not at all necessary to follow the melting body with your eyes. It is enough to look at the mercury column of the thermometer. Until melting has begun, the body temperature rises (Fig. 4.5). As soon as melting starts, the temperature rise stops and the temperature will remain unchanged until the melting process is complete.

Rice. 4.5

Like the transformation of a liquid into vapor, the transformation of a solid into a liquid requires heat. The heat required for this is called the latent heat of fusion. For example, melting one kilogram of ice requires 80 kcal.

Ice is one of the bodies with a high heat of fusion. Melting ice requires, for example, 10 times more energy than melting the same mass of lead. Of course, we are talking about the melting itself, we are not saying here that before the melting of lead begins, it must be heated to + 327 ° C. Due to the high heat of melting ice, the melting of snow slows down. Imagine that the heat of melting would be 10 times less. Then spring floods would bring unimaginable disasters every year.

So, the heat of melting of ice is great, but it is also small if compared with the specific heat of vaporization of 540 kcal/kg (seven times less). However, this difference is quite natural. When converting a liquid into vapor, we must tear the molecules one from the other, and when melting, we only have to destroy the order in the arrangement of the molecules, leaving them at almost the same distances. It is clear that less work is required in the second case.

The presence of a certain melting point is an important feature of crystalline substances. It is on this basis that they are easy to distinguish from other solids, called amorphous or glasses. Glasses are found among both inorganic and organic substances. Window panes are usually made from sodium and calcium silicates; on the desk often put organic glass (it is also called plexiglass).

Amorphous substances, in contrast to crystals, do not have a definite melting point. Glass does not melt, but softens. When heated, a piece of glass first becomes soft from hard, it can be easily bent or stretched; at a higher temperature, the piece begins to change its shape under the influence of its own gravity. As it heats up, the thick viscous mass of glass takes the shape of the vessel in which it lies. This mass is at first thick, like honey, then like sour cream, and, finally, it becomes almost as low-viscosity liquid as water. With all our desire, we cannot indicate here a specific temperature for the transition of a solid to a liquid. The reasons for this lie in the fundamental difference between the structure of glass and the structure of crystalline bodies. As mentioned above, the atoms in amorphous bodies arranged randomly. Glasses in structure resemble liquids. Even in solid glass, the molecules are arranged randomly. This means that an increase in the temperature of glass only increases the range of vibrations of its molecules, giving them gradually more and more freedom of movement. Therefore, the glass softens gradually and does not show a sharp "solid" - "liquid" transition, which is characteristic of the transition from the arrangement of molecules in a strict order to a random arrangement.

When it came to the boiling curve, we said that liquid and vapor can, albeit in an unstable state, live in foreign regions - vapor can be supercooled and transferred to the left of the boiling curve, liquid can be superheated and pulled to the right of this curve.

Are similar phenomena possible in the case of a crystal with a liquid? It turns out that the analogy here is incomplete.

If you heat the crystal, it will begin to melt at its melting point. The crystal cannot be overheated. On the contrary, by cooling the liquid, it is possible, if certain measures are taken, to “slip through” the melting point relatively easily. In some liquids, large subcoolings can be achieved. There are even liquids that are easy to supercool, but difficult to make crystallize. As such a liquid cools, it becomes more and more viscous and finally solidifies without crystallizing. Such is glass.

You can also recool water. The mist droplets may not freeze even when severe frosts. If a crystal of a substance, a seed, is thrown into a supercooled liquid, then crystallization will immediately begin.

Finally, in many cases delayed crystallization may be initiated by a shake or other random events. It is known, for example, that crystalline glycerol was first obtained during transportation through railway. Glasses after a long standing may begin to crystallize (devitrify, or "collapse", as they say in technology).

How to grow a crystal

Almost any substance can give crystals under certain conditions. Crystals can be obtained from a solution or melt of a given substance, as well as from its vapor (for example, black diamond-shaped crystals of iodine easily precipitate from its vapor at normal pressure without an intermediate transition to a liquid state).

Start dissolving table salt or sugar in the water. At room temperature (20°C), you will be able to dissolve only 70 g of salt in a faceted glass. Further additions of salt will not dissolve and will settle at the bottom in the form of sediment. A solution in which no further dissolution occurs is called saturated. .If you change the temperature, then the degree of solubility of the substance will also change. Everyone is well aware that hot water dissolves most substances much easier than cold water.

Imagine now - that you have prepared a saturated solution of, say, sugar at a temperature of 30 ° C and begin to cool it to 20 ° C. At 30°C, you were able to dissolve 223 g of sugar in 100 g of water, at 20°C, 205 g dissolves. Then, when cooled from 30 to 20°C, 18 g will be "extra" and, as they say, will fall out of solution. So, one of the possible ways to obtain crystals is to cool the saturated solution.

You can do it differently. Prepare a saturated salt solution and leave it in an open glass. After a while, you will find the appearance of crystals. Why did they form? Careful observation will show that simultaneously with the formation of crystals, another change occurred - the amount of water decreased. The water evaporated, and the "extra" substance appeared in the solution. So another possible way the formation of crystals is the evaporation of a solution.

How does crystals form from solution?

We said that the crystals "fall out" of the solution; Is it necessary to understand this in such a way that there was no crystal for a week, and in one moment it suddenly appeared at once? No, this is not the case: the crystals grow. It is not possible, of course, to detect the very initial moments of growth with the eye. At first, a few of the randomly moving molecules or atoms of the solute assemble in the approximate order needed to form the crystal lattice. Such a group of atoms or molecules is called a nucleus.

Experience shows that nuclei are more often formed in the presence of any extraneous minute dust particles in the solution. The fastest and easiest crystallization begins when a small seed crystal is placed in a saturated solution. In this case, the isolation of a solid from the solution will not consist in the formation of new crystals, but in the growth of the seed.

The growth of the embryo does not, of course, differ from the growth of the seed. The meaning of using a seed is that it "pulls" the released substance onto itself and thus prevents the simultaneous formation a large number embryos. If many nuclei are formed, then they will interfere with each other during growth and will not allow us to obtain large crystals.

How are the portions of atoms or molecules released from the solution distributed on the surface of the nucleus?

Experience shows that the growth of a nucleus or a seed consists, as it were, in moving the faces parallel to themselves in a direction perpendicular to the face. In this case, the angles between the faces remain constant (we already know that the constancy of angles is the most important feature of a crystal, which follows from its lattice structure).

On fig. 4.6 the outlines of three crystals of the same substance that occur during their growth are given. Similar patterns can be observed under a microscope. In the case shown on the left, the number of faces is conserved during growth. The middle drawing gives an example of a new face appearing (upper right) and disappearing again.

Rice. 4.6

It is very important to note that the growth rate of the faces, i.e., the speed of their movement parallel to themselves, is not the same for different faces. In this case, exactly those faces that move the fastest, for example, the lower left face in the middle figure, "overgrow" (disappear). On the contrary, slowly growing faces are the widest, as they say, the most developed.

This is especially clear in the last figure. The shapeless fragment acquires the same shape as other crystals precisely because of the growth rate anisotropy. Well-defined facets develop at the expense of others most strongly and give the crystal a form characteristic of all samples of this substance.

Very beautiful transitional forms are observed when a ball is taken as a seed, and the solution is alternately slightly cooled and heated. When heated, the solution becomes unsaturated, and the seed is partially dissolved. Cooling leads to saturation of the solution and growth of the seed. But the molecules settle in a different way, as if giving preference to certain places. The substance is thus transferred from one place of the ball to another.

First, small circle-shaped faces appear on the surface of the ball. The circles gradually increase and, touching each other, merge along straight edges. The ball turns into a polyhedron. Then some faces overtake others, some of the faces overgrow, and the crystal acquires its characteristic shape (Fig. 4.7).

Rice. 4.7

When observing the growth of crystals, the main feature of growth is striking - the parallel movement of the faces. It turns out that the released substance builds up the face in layers: until one layer is completed, the next one does not begin to build.

On fig. 4.8 shows the "unfinished" packing of atoms. In which of the positions indicated by letters will the new atom be most firmly held, attached to the crystal? No doubt in A, since here he experiences the attraction of neighbors from three sides, while in B - from two, and in C - only from one side. Therefore, the column is first completed, then the entire plane, and only then the laying of a new plane begins.

Rice. 4.8

In a number of cases, crystals are formed from a molten mass - from a melt. In nature, this happens on an enormous scale: basalts, granites and many other rocks arose from fiery magma.

Let's start heating some crystalline substance, for example, rock salt. Up to 804°C, rock salt crystals will change little: they expand only slightly, and the substance remains solid. A temperature meter placed in a vessel with a substance shows a continuous increase in temperature when heated. At 804°C, we will immediately discover two new, interconnected phenomena: the substance will begin to melt, and the rise in temperature will stop. Until all matter turns into a liquid,; the temperature will not change; a further rise in temperature is already heating the liquid. All crystalline substances have a certain melting point. Ice melts at 0°C, iron melts at 1527°C, mercury melts at -39°C, etc.

As we already know, in each crystal the atoms or molecules of a substance form an ordered G packing and make small vibrations around their average positions. As the body heats up, the speed of the oscillating particles increases along with the amplitude of the oscillations. This increase in the speed of particles with increasing temperature is one of the basic laws of nature, which applies to matter in any state - solid, liquid or gaseous.

When a certain, sufficiently high temperature of the crystal is reached, the oscillations of its particles become so energetic that an accurate arrangement of the particles becomes impossible - the crystal melts. With the onset of melting, the heat supplied is no longer used to increase the particle velocity, but to destroy the crystal lattice. Therefore, the rise in temperature is suspended. Subsequent heating is an increase in the velocity of the liquid particles.

In the case of crystallization from a melt that interests us, the above phenomena are observed in the reverse order: as the liquid cools, its particles slow down their chaotic motion; when a certain, sufficiently low temperature is reached, the velocity of the particles is already so low that some of them, under the action of attractive forces, begin to attach themselves to one another, forming crystalline nuclei. Until all the substance crystallizes, the temperature remains constant. This temperature is generally the same as the melting point.

If special measures are not taken, then crystallization from the melt will begin immediately in many places. Crystals will grow in the form of regular polyhedrons characteristic of them in exactly the same way as we described above. However, free growth does not last long: growing, the crystals collide with each other, growth stops at the points of contact, and the solidified body acquires a granular structure. Each grain is a separate crystal, which failed to take its correct form.

Depending on many conditions, and above all on the rate of cooling, a solid body may have more or less large grains: the slower the cooling, the larger the grains. The grain sizes of crystalline bodies range from a millionth of a centimeter to several millimeters. In most cases, the granular crystalline structure can be observed under a microscope. Solids usually have just such a fine-grained structure.

For technology, the process of solidification of metals is of great interest. The events that occur during casting and during the solidification of metal in molds have been studied by physicists in great detail.

For the most part, during solidification, tree-like single crystals grow, which are called dendrites. In other cases, the dendrites are oriented randomly, in other cases, they are parallel to each other.

On fig. 4.9 shows the stages of growth of one dendrite. With this behavior, a dendrite can overgrow before it meets another similar one. Then we will not find dendrites in the casting. Events can also develop differently: dendrites can meet and grow into each other (branches of one in the gaps between the branches of another) while they are still "young".

Rice. 4.9

In this way, castings may arise whose grains (shown in Fig. 2.22) have the most different structure. And the properties of metals significantly depend on the nature of this structure. It is possible to control the behavior of the metal during solidification by changing the cooling rate and the heat removal system.

Now let's talk about how to grow a large single crystal. It is clear that measures must be taken to ensure that the crystal grows from one place. And if several crystals have already begun to grow, then in any case it is necessary to make sure that the growth conditions are favorable for only one of them.

Here, for example, is how they proceed when growing crystals of low-melting metals. The metal is melted in a glass test tube with a drawn end. A test tube suspended by a thread inside a vertical cylindrical furnace is slowly lowered down. The drawn end gradually exits the furnace and cools. Crystallization begins. At first, several crystals form, but those that grow sideways rest against the wall of the test tube and their growth slows down. Only the crystal that grows along the axis of the test tube, i.e., deep into the melt, will be in favorable conditions. As the test tube is lowered, new portions of the melt, falling into the region of low temperatures, will "feed" this single crystal. Therefore, of all the crystals, he alone survives; as the tube is lowered, it continues to grow along its axis. In the end, all the molten metal solidifies in the form of a single crystal.

The same idea underlies the growth of refractory ruby ​​crystals. A fine powder of the substance is squirted through the flame. At the same time, the powders melt; tiny drops fall on a refractory support of a very small area, forming many crystals. As the drops fall further onto the stand, all the crystals grow, but again, only the one that is in the most favorable position for "receiving" the falling drops grows.

What are large crystals for?

Industry and science often need large single crystals. Great importance for technology, they have crystals of Rochelle salt and quartz, which have the remarkable property of converting mechanical actions (for example, pressure) into electrical voltage.

The optical industry needs large crystals of calcite, rock salt, fluorite, etc.

The watch industry needs crystals of rubies, sapphires and some other precious stones. The fact is that individual moving parts of ordinary watches make up to 20,000 vibrations per hour. Such a high load places unusually high demands on the quality of the axle tips and bearings. Abrasion will be the smallest when a ruby ​​or sapphire serves as a bearing for the tip of an axle with a diameter of 0.07-0.15 mm. Artificial crystals of these substances are very durable and are very little abraded by steel. It's great that artificial stones are thus better than the same natural stones.

However, the growth of single crystals of semiconductors - silicon and germanium - is of the greatest importance for industry.

The influence of pressure on the melting point

If the pressure is changed, the melting point will also change. We met with the same regularity when we talked about boiling. The more pressure; the higher the boiling point. As a rule, this is also true for melting. However, there are a small number of substances that behave anomalously: their melting point decreases with increasing pressure.

The fact is that the vast majority of solids are denser than their liquids. The exception to this dravil is precisely those substances whose melting point does not change quite normally with a change in pressure, for example, water. Ice is lighter than water, and the melting point of ice decreases as pressure increases.

Compression promotes the formation of a denser state. If a solid is denser than a liquid, then compression helps solidify and prevents melting. But if melting is hampered by compression, then this means that the substance remains solid, whereas earlier at this temperature it would have already melted, i.e., with increasing pressure, the melting point increases. AT abnormal case a liquid is denser than a solid, and pressure helps the formation of the liquid, i.e., lowers the melting point.

The effect of pressure on the melting point is much less than that of boiling. An increase in pressure by more than 100 kgf / cm 2 lowers the melting point of ice by 1°C.

Why do skates glide only on ice, but not on equally smooth parquet? Apparently, the only explanation is the formation of water, which lubricates the skate. To understand the contradiction that has arisen, we need to remember the following: blunt skates slide very poorly on ice. Skates need to be sharpened to cut ice. In this case, only the tip of the edge of the ridge presses on the ice. The pressure on the ice reaches tens of thousands of atmospheres, the ice still melts.

Evaporation of solids

When they say "a substance evaporates", they usually mean that a liquid evaporates. But solids can also evaporate. Sometimes the evaporation of solids is called sublimation.

The evaporating solid is, for example, naphthalene. Naphthalene melts at 80°C and evaporates at room temperature. It is this property of naphthalene that allows it to be used to exterminate moths.

A fur coat covered with naphthalene is saturated with naphthalene vapor and creates an atmosphere that moths cannot stand. Any smelling solid sublimes to a large extent. After all, the smell is created by molecules that have broken away from the substance and reached our nose. However, there are more frequent cases where the substance is sublimated to an insignificant degree, sometimes to a degree that cannot be detected even by very careful research. In principle, any solid substance (precisely any, even iron or copper) evaporates. If we do not detect sublimations, this only means that the density of the saturating vapor is very low.

It can be seen that a number of substances that have a pungent odor at room temperature lose it at low temperature.

The density of saturated vapor in equilibrium with a solid increases rapidly with increasing temperature. We illustrated this behavior with the curve for ice shown in Fig. 4.10. True, the ice does not smell ...

Rice. 4.10

In most cases, it is impossible to significantly increase the density of saturated vapor of a solid for a simple reason - the substance will melt earlier.

The ice also evaporates. This is well known to housewives who hang out wet laundry to dry in cold weather. The water first freezes, and then the ice evaporates, and the laundry turns out to be dry.

triple point

So, there are conditions under which vapor, liquid and crystal can exist in pairs in equilibrium. Can all three states be in equilibrium? Such a point on the pressure-temperature diagram exists, it is called triple. Where is she?

If you place water with floating ice in a closed vessel at zero degrees, then in free space water (and "ice") vapors will begin to flow. At a vapor pressure of 4.6 mm Hg. Art. Evaporation will stop and saturation will begin. Now the three phases - ice, water and steam - will be in equilibrium. This is the triple point.

The relationship between the various states is clearly and clearly shown by the diagram for water shown in fig. 4.11.

Rice. 4.11

Such a diagram can be constructed for any body.

The curves in the figure are familiar to us - these are equilibrium curves between ice and steam, ice and water, water and steam. As usual, pressure is plotted vertically, and temperature is plotted horizontally.

The three curves intersect at the triple point and divide the diagram into three areas - the living spaces of ice, water and water vapor.

The state diagram is a concise reference. Its purpose is to answer the question of what state of the body is stable at such and such a pressure and such and such a temperature.

If water or steam is placed in the conditions of the "left region", they will become ice. If a liquid or a solid body is introduced into the "lower region", then steam will be obtained. In the "right region" the vapor will condense and the ice will melt.

The diagram of the existence of phases allows you to immediately answer what happens to the substance when heated or when compressed. Heating at a constant pressure is shown as a horizontal line in the diagram. A dot moves along this line from left to right, representing the state of the body.

The figure shows two such lines, one of them is heating at normal pressure. The line lies above the triple point. Therefore, it will cross first the melting curve, and then, outside the drawing, the evaporation curve. Ice at normal pressure will melt at 0°C, and the resulting water will boil at 100°C.

The situation will be different for ice heated at very low pressure, say just below 5 mm Hg. Art. The heating process is represented by a line below the triple point. The melting and boiling curves do not intersect with this line. At such a slight pressure, heating will lead to a direct transition of ice into steam.

On fig. 4.12, the same diagram shows what an interesting phenomenon will occur when water vapor is compressed in the state marked with a cross in the figure. The steam will first turn into ice and then melt. The figure allows you to immediately tell at what pressure the growth of the crystal will begin and when the melting will occur.

Rice. 4.12

State diagrams of all substances are similar to each other. Large, from an everyday point of view, differences arise due to the fact that the location of the triple point on the diagram can be very different for different substances.

After all, we exist near "normal conditions", that is, primarily at a pressure close to one atmosphere. How the triple point of matter is located in relation to the line of normal pressure is very important for us.

If the pressure at the triple point is less than atmospheric, then for us, living in "normal" conditions, the substance is melting. When the temperature rises, it first turns into a liquid, and then boils.

In the opposite case - when the pressure at the triple point is higher than atmospheric - we will not see liquid when heated, the solid will directly turn into vapor. This is how "dry ice" behaves, which is very convenient for ice cream sellers. Blocks of ice cream can be shifted with pieces of "dry ice" and not be afraid that the ice cream will become wet. "Dry ice" is solid carbon dioxide CO 2 . The triple point of this substance lies at 73 atm. Therefore, when solid CO 2 is heated, the point depicting its state moves horizontally, crossing only the evaporation curve of the solid (same as for regular ice at a pressure of about 5 mm Hg. Art.).

We have already told the reader how one degree of temperature is determined on the Kelvin scale, or, as the SI system now requires, one kelvin. However, it was about the principle of determining the temperature. Not all metrology institutes have ideal gas thermometers. Therefore, the temperature scale is built with the help of equilibrium points fixed by nature between different states of matter.

The triple point of water plays a special role in this. The degree Kelvin is now defined as 273.16th of the thermodynamic temperature of the triple point of water. The triple point of oxygen is taken equal to 54.361 K. The solidification temperature of gold is set to 1337.58 K. Using these reference points, any thermometer can be accurately calibrated.

The same atoms, but ... different crystals

The matte black soft graphite we write with and the brilliant, transparent, hard, glass-cutting diamond are built from the same carbon atoms. Why are the properties of these two identical substances so different?

Recall the lattice of layered graphite, each atom of which has three nearest neighbors, and the lattice of diamond, whose atom has four nearest neighbors. This example clearly shows that the properties of crystals are determined by the mutual arrangement of atoms. Graphite is used to make refractory crucibles that can withstand temperatures up to two to three thousand degrees, and diamond burns at temperatures above 700 ° C; the density of diamond is 3.5, and that of graphite is 2.3; graphite conducts electricity, diamond - does not conduct, etc.

It is not only carbon that has this feature of producing different crystals. Almost every chemical element, and not only an element, but any chemical substance, can exist in several varieties. Six varieties of ice, nine varieties of sulfur, four varieties of iron are known.

When discussing the state diagram, we did not talk about different types crystals and drew a single area of ​​the solid. And this area for very many substances is divided into sections, each of which corresponds to a certain "grade" of a solid body or, as they say, a certain solid phase (a certain crystalline modification).

Each crystalline phase has its own region of stable state, limited by a certain range of pressures and temperatures. The laws of transformation of one crystalline variety into another are the same as the laws of melting and evaporation.

For each pressure, you can specify the temperature at which both types of crystals will peacefully coexist. If the temperature is increased, a crystal of one kind will turn into a crystal of the second kind. If the temperature is lowered, the reverse transformation will occur.

In order for red sulfur to turn yellow at normal pressure, a temperature below 110 ° C is needed. Above this temperature, up to the melting point, the arrangement of atoms characteristic of red sulfur is stable. The temperature drops, the vibrations of the atoms decrease, and, starting from 110 ° C, nature finds a more convenient arrangement of atoms. There is a transformation of one crystal into another.

six different ice no one came up with a name. So they say: ice one, ice two, ...., ice seven. How about seven, if there are only six varieties? The fact is that ice four was not detected during repeated experiments.

If water is compressed at a temperature of about zero, then at a pressure of about 2000 atm ice five is formed, and at a pressure of about 6000 atm ice six is ​​formed.

Ice two and ice three are stable at temperatures below zero degrees.

Ice seven - hot ice; it comes from compression hot water up to pressures of about 20,000 atm.

All ice, except ordinary ice, is heavier than water. Ice produced under normal conditions behaves anomalously; on the contrary, ice obtained under conditions different from the norm behaves normally.

We say that each crystalline modification is characterized by a certain area of ​​existence. But if so, how do graphite and diamond exist under the same conditions?

Such "lawlessness" in the world of crystals is very common. The ability to live in "foreign" conditions for crystals is almost the rule. If in order to transfer a vapor or a liquid to other areas of existence, one has to resort to various tricks, then a crystal, on the contrary, can almost never be forced to remain within the boundaries assigned to it by nature.

Overheating and supercooling of crystals are explained by the difficulty of converting one order into another under conditions of extreme crowding. Yellow sulfur should turn red at 95.5°C. With more or less rapid heating, we will "skip" this transformation point and bring the temperature up to the melting point of sulfur 113°C.

The true transformation temperature is easiest to detect when the crystals come into contact. If they are closely placed one on top of the other and kept at 96°C, then the yellow will be eaten by the red, and at 95°C the yellow will absorb the red. In contrast to the "crystal-liquid" transition, the "crystal-crystal" transformations are usually delayed both during supercooling and overheating.

In some cases, we are dealing with such states of matter, which would be supposed to live at completely different temperatures.

White tin should turn gray when the temperature drops to +13°C. We usually deal with white tin and know that nothing is done with it in winter. It perfectly withstands hypothermia of 20-30 degrees. However, in severe winter conditions, white tin turns into gray. Ignorance of this fact was one of the circumstances that ruined Scott's expedition to South Pole(1912). The liquid fuel taken by the expedition was in vessels brazed with tin. In great colds, white tin turned into a gray powder - the vessels were unsoldered; and the fuel spilled out. No wonder the appearance of gray spots on white tin is called tin plague.

Just as in the case of sulfur, white tin can be turned into gray at a temperature just below 13 ° C; if only a tiny grain of the gray variety falls on a pewter object.

The existence of several varieties of the same substance and delays in their mutual transformations are of great importance for technology.

At room temperature, iron atoms form a body-centered cubic lattice in which the atoms occupy positions at the vertices and in the center of the cube. Each atom has 8 neighbors. At high temperatures, iron atoms form a denser "packing" - each atom has 12 neighbors. Iron with 8 neighbors is soft, iron with 12 neighbors is hard. It turns out that it is possible to obtain iron of the second type at room temperature. This method - hardening - is widely used in metallurgy.

Hardening is carried out very simply - a metal object is red-hot, and then thrown into water or oil. Cooling occurs so rapidly that the transformation of the structure, which is stable at high temperature, does not have time to occur. Thus, a high-temperature structure will exist indefinitely under conditions unusual for it: recrystallization into a stable structure proceeds so slowly that it is practically imperceptible.

Speaking about the hardening of iron, we were not entirely accurate. Steel is tempered, i.e. iron containing fractions of a percent of carbon. The presence of very small carbon impurities delays the transformation of hard iron into soft and allows hardening. As for completely pure iron, it is not possible to harden it - the transformation of the structure has time to occur even with the most abrupt cooling.

Depending on the type of state diagram, by changing the pressure or temperature, certain transformations are achieved.

Many crystal-to-crystal transformations are observed with a change in pressure alone. In this way, black phosphorus was obtained.

Rice. 4.13

It was possible to turn graphite into diamond only by using both high temperature, and great pressure. On fig. 4.13 shows the state diagram of carbon. At pressures below ten thousand atmospheres and at temperatures below 4000 K, graphite is a stable modification. Thus, the diamond lives in "foreign" conditions, so it can be easily turned into graphite. But the inverse problem is of practical interest. It is not possible to carry out the transformation of graphite into diamond only by increasing the pressure. The phase transformation in the solid state apparently proceeds too slowly. The appearance of the state diagram suggests the correct solution: increase the pressure and heat at the same time. Then we get (right corner of the diagram) molten carbon. Cooling it down high pressure, we have to get into the area of ​​the diamond.

The practical possibility of such a process was proved in 1955, and at present the problem is considered to be technically solved.

Amazing Liquid

If you lower the body temperature, then sooner or later it will harden and acquire a crystalline structure. It does not matter at what pressure the cooling occurs. This circumstance seems quite natural and understandable from the point of view of the laws of physics, with which we have already become acquainted. Indeed, by lowering the temperature, we reduce the intensity of thermal motion. When the movement of molecules becomes so weak that it no longer interferes with the forces of interaction between them, the molecules line up in a neat order - they form a crystal. Further cooling will take away from the molecules all the energy of their movement, and at absolute zero the substance must exist in the form of resting molecules arranged in a regular lattice.

Experience shows that all substances behave in this way. All, except for one and only: such a "freak" is helium.

We have already given the reader some information about helium. Helium holds the record for its critical temperature. No substance has a critical temperature lower than 4.3 K. However, this record in itself does not mean anything surprising. Another thing is striking: by cooling helium below the critical temperature, reaching almost absolute zero, we will not get solid helium. Helium remains liquid even at absolute zero.

The behavior of helium is completely inexplicable from the point of view of the laws of motion we have outlined and is one of the signs of the limited validity of such laws of nature, which seemed to be universal.

If the body is liquid, then its atoms are in motion. But after all, having cooled the body to absolute zero, we took away all the energy of movement from it. We have to admit that helium has such an energy of motion that cannot be taken away. This conclusion is incompatible with the mechanics we have been dealing with so far. According to this mechanics we have studied, the movement of a body can always be slowed down to a complete stop by taking away all its kinetic energy; in the same way, it is possible to stop the movement of molecules by taking away their energy when they collide with the walls of a cooled vessel. For helium, such mechanics is clearly not suitable.

The "strange" behavior of helium is an indication of a fact of great importance. We first met with the impossibility of applying in the world of atoms the basic laws of mechanics, established by a direct study of the motion of visible bodies, laws, which seemed to be the unshakable foundation of physics.

The fact that helium "refuses" to crystallize at absolute zero cannot be reconciled in any way with the mechanics we have studied so far. The contradiction with which we met for the first time - the disobedience of the world of atoms to the laws of mechanics - is only the first link in the chain of even sharper and sharper contradictions in physics.

These contradictions lead to the need to revise the foundations of the mechanics of the atomic world. This revision is very profound and leads to a change in our entire understanding of nature.

The need for a radical revision of the mechanics of the atomic world does not mean that we should put an end to the laws of mechanics we have studied. It would be unfair to force the reader to learn unnecessary things. The old mechanics is completely valid in the world of large bodies. Already this is enough to treat the relevant chapters of physics with full respect. However, it is also important that a number of laws of the "old" mechanics pass into the "new" mechanics. This includes, in particular, the law of conservation of energy.

The presence of "irremovable" energy at absolute zero is not a special property of helium. Turns out; "zero" energy is present in all substances.

Only in helium this energy is enough to prevent the atoms from forming the correct crystal lattice.

It is not necessary to think that helium cannot be in a crystalline state. For the crystallization of helium, it is only necessary to increase the pressure to about 25 atm. Cooling carried out at a higher pressure will lead to the formation of solid crystalline helium with quite ordinary properties. Helium forms a face-centered cubic lattice.

On fig. 4.14 shows a diagram of the state of helium. It differs sharply from the diagrams of all other substances in the absence of a triple point. The melting and boiling curves do not intersect.

Rice. 4.14

And this unique state diagram has one more feature: there are two different helium liquids. What is their difference - you will learn a little later.