Dependence of boiling temperature on pressure examples. Boiling of liquids

From the above considerations it is clear 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. Thus, 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, water boiling in hermetically sealed vessels - autoclaves (Fig. 6.11) also occurs at elevated pressure. Therefore, the boiling point is significantly higher than 100 °C. Autoclaves are used to sterilize surgical instruments, dressings, etc.

And vice versa, by reducing external pressure, we thereby lower the boiling point. Under the bell of an air pump, you can make water boil at room temperature (Fig. 6.12). As you climb mountains, the atmospheric pressure decreases, therefore 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, for example, under these conditions.

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

Differences in boiling points of liquids

Each liquid has its own boiling point. The difference in boiling points of liquids is determined by the difference in the pressure of their saturated vapors at the same temperature. For example, ether vapors already at room temperature have a pressure greater than half atmospheric. Therefore, in order for the ether vapor pressure to become equal to atmospheric pressure, a slight increase in temperature (up to 35 ° C) is necessary. In mercury, saturated vapors have room temperature very little pressure. The pressure of mercury vapor becomes equal to atmospheric pressure 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 boiling points of substances is widely used in technology, for example, in the separation of petroleum products. When oil is heated, its most valuable, volatile parts (gasoline) evaporate first, 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 change liquid into vapor? Most likely yes! Isn't 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 from outside. Of course, heat itself can be transferred to the liquid from surrounding bodies. So, the water in the glass evaporates, but the temperature of the water, slightly lower than the ambient temperature, remains unchanged. Heat is transferred from air to water until all the water has evaporated.

To maintain the boiling of water (or other liquid), 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 increase, but a certain amount of steam is produced every second.

Thus, to convert a liquid into vapor by evaporation or by boiling, an input 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 this liquid.

What is the energy supplied to the body spent on? First of all, to increase its internal energy during the transition from liquid state into gaseous: after all, in this case the volume of the 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, as the volume of a substance increases, work is done against external pressure forces. This part of the heat of vaporization at room temperature is usually several 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 liquid vaporization; T- mass of liquid; Q n- its heat of vaporization. The SI unit of specific heat of vaporization is 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.

Using the phenomenon of cooling a liquid as it evaporates; dependence of the boiling point of water on pressure.

During vaporization, a substance passes from a liquid state to a gaseous state (steam). There are two types of vaporization: evaporation and boiling.

Evaporation is vaporization occurring from the free surface of a liquid.

How does evaporation occur? We know that the molecules of any liquid are in continuous and random motion, some of them moving faster, others slower. They are prevented from flying out by the forces of attraction towards each other. If, however, there is a molecule with a sufficiently high kinetic energy at the surface of the liquid, then it will be able to overcome the forces of intermolecular attraction and fly out of the liquid. The same thing will be repeated with another fast molecule, with the second, third, etc. Flying out, these molecules form vapor above the liquid. The formation of this steam is evaporation.

Since the fastest molecules fly out of a liquid during evaporation, the average kinetic energy There are fewer and fewer molecules remaining in the liquid. As a result of this the temperature of the evaporating liquid decreases: The liquid is cooled. This is why, in particular, a person in wet clothes feels colder than in dry clothes (especially in the wind).

At the same time, everyone knows that if you pour water into a glass and leave it on the table, then, despite evaporation, it will not cool continuously, becoming colder and colder until it freezes. What's stopping this? The answer is very simple: heat exchange between water and the warm air surrounding the glass.

Cooling of a liquid during evaporation is more noticeable in the case when evaporation occurs quickly enough (so that the liquid does not have time to restore its temperature due to heat exchange with environment). Volatile liquids with weak intermolecular attractive forces, such as ether, alcohol, and gasoline, evaporate quickly. If you drop such a liquid on your hand, you will feel cold. Evaporating from the surface of the hand, such a liquid will cool and take away some heat from it.

Rapidly evaporating substances are found wide application in technology. For example, in space technology, descent vehicles are coated with such substances. When passing through the planet's atmosphere, the body of the apparatus heats up as a result of friction, and the substance covering it begins to evaporate. As it evaporates, it cools the spacecraft, thereby saving it from overheating.

Cooling of water during its evaporation is also used in instruments used to measure air humidity - psychrometers(from the Greek “psychros” - cold). The psychrometer consists of two thermometers. One of them (dry) shows the air temperature, and the other (the reservoir of which is tied with cambric dipped in water) shows more low temperature, due to the intensity of evaporation from wet cambric. The drier the air whose humidity is measured, the greater the evaporation and therefore the lower the wet-bulb reading. And vice versa, the higher the air humidity, the less intense evaporation occurs and therefore the higher the temperature this thermometer shows. Based on the readings of dry and humidified thermometers, air humidity, expressed as a percentage, is determined using a special (psychrometric) table. The highest humidity is 100% (at this air humidity, dew appears on objects). For humans, the most favorable humidity is considered to be between 40 and 60%.

With the help of simple experiments it is easy to establish that the rate of evaporation increases with increasing temperature of the liquid, as well as with increasing area of ​​its free surface and in the presence of wind.

Why does liquid evaporate faster when there is wind? The fact is that simultaneously with evaporation on the surface of the liquid, the reverse process also occurs - condensation. Condensation occurs due to the fact that some of the vapor molecules, moving randomly over the liquid, return to it again. The wind carries away the molecules that fly out of the liquid and does not allow them to return back.

Condensation can also occur when the vapor is not in contact with the liquid. It is condensation, for example, that explains the formation of clouds: molecules of water vapor rising above the ground in the colder layers of the atmosphere are grouped into tiny droplets of water, the accumulations of which constitute clouds. The condensation of water vapor in the atmosphere also results in rain and dew.

Dependence of boiling temperature on pressure

The boiling point of water is 100°C; one might think that this is an inherent property of water, that water, no matter where and under what conditions it is, will always boil at 100°C.

But this is not so, and residents of high mountain villages are well aware of this.

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

What's the matter? What physical factor interferes with the boiling phenomenon? What is the significance of altitude above sea level?

This physical factor is the pressure acting on the surface of the liquid. You don't need to climb to the top of a mountain to verify the truth of what has been said.

By placing heated water under a bell and pumping or pumping out air from there, you can make sure that the boiling point rises as the pressure increases and falls as it decreases.

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 refresh yourself 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 cause the water to 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 pumping continues, the pressure will also drop. 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 layer of water near a metal surface drops greatly and the water in this layer boils, i.e., numerous steam-filled bubbles appear in it. This phenomenon is called cavitation (from Latin word cavitas - cavity).

By reducing the pressure, we lower the boiling point. And by increasing it? A graph like ours answers this question. A pressure of 15 atm can delay the boiling of water, it will begin only at 200°C, and a pressure of 80 atm will cause water to boil only at 300°C.

So, a certain external pressure corresponds to a certain boiling point. But this statement can be “turned around” by 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 a curve of vapor pressure as a function of temperature.

The numbers plotted on the boiling point graph (or on the vapor pressure graph) show that vapor pressure changes very sharply with temperature. At 0°C (i.e. 273 K) the vapor pressure is 4.6 mmHg. Art., at 100°C (373 K) it is equal to 760 mm Hg. Art., i.e. increases 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. almost up to 60 atm, i.e. approximately 10,000 times.

Therefore, on the contrary, the boiling point changes with pressure rather slowly. When the pressure doubles 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 doubles 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 (condensation) of steam into water.

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

Both during boiling and during condensation, the point will not move from the curve until the conversion of steam into water or water into steam is complete. This can also be formulated this way: under the conditions of our curve and only under these conditions, the coexistence of liquid and vapor is possible. If you do not add or remove heat, then the amounts of steam and liquid in a closed vessel will remain unchanged. Such vapor and liquid are said to be in equilibrium, and vapor that is in equilibrium with its liquid is called saturated.

The boiling and condensation curve, 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. Left and up (to high temperatures and lower pressures) there is a region of stable state of steam. To the right and down is the region of the stable state of the liquid.

The vapor-liquid equilibrium curve, i.e. the curve of the dependence of boiling point on pressure or, which 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 the vapor pressure always increases rapidly with increasing temperature.

We have already used the words “gas” and “steam” many times. These two words are pretty equal. We can say: water gas is water vapor, oxygen gas is oxygen liquid vapor. Nevertheless, a certain habit has developed when using 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 elasticity at ordinary temperatures is higher than atmospheric pressure. On the contrary, we talk about vapor when, at room temperature and atmospheric pressure, the substance is more stable in the form of a liquid.

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

The liquid not only evaporates. At a certain temperature it boils.
Dependence of saturated vapor pressure on 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 temperature increases, pressure increases. 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, like that of an ideal gas at constant volume. With increasing temperature, the pressure of real saturated vapor increases faster than the pressure of an ideal gas ( Fig.11.1, part of the curve AB). This becomes obvious if we draw isochores of an ideal gas through the points A And IN(dashed lines). Why is this happening?

When a liquid is heated in a closed container, some of the liquid turns into steam. 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 steam in a closed vessel changes (or when the volume changes at a constant temperature), the mass of the steam changes. The liquid partially turns into vapor, or, on the contrary, the vapor partially condenses. WITH ideal gas nothing like that happens.
When all the liquid has evaporated, the steam will cease to be saturated upon further heating and its pressure at a constant volume will increase in direct proportion absolute temperature(cm. Fig.11.1, part 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 are formed throughout the entire volume of the liquid, which float to the surface. The boiling point of the liquid remains constant. This happens because all the energy supplied to the liquid is spent converting it into vapor. Under what conditions does boiling begin?
A liquid always contains dissolved gases, released at the bottom and walls of the vessel, as well as on dust particles suspended in the liquid, which are centers of vaporization. The liquid vapors inside the bubbles are saturated. With increasing temperature the pressure saturated vapors increases and the bubbles increase in size. Under the influence of buoyant force they float upward. If the upper layers of the liquid have a lower temperature, then vapor condensation occurs in bubbles in these layers. The pressure drops rapidly and the bubbles collapse. The collapse occurs so quickly that the walls of the bubble collide, producing something like an explosion. Many such micro-explosions create a characteristic noise. When the liquid warms up enough, the bubbles will 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 it boils.
The dependence of saturated 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 below the boiling point, and only from the surface of the liquid; during boiling, vapor formation occurs throughout the entire volume of the liquid.
Boiling begins at the temperature at which the saturated vapor pressure in the bubbles is equal to the pressure in the liquid.
The greater the external pressure, the higher the boiling point. Thus, 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 in hermetically sealed vessels - autoclaves ( Fig.11.2) boiling of water also occurs at elevated pressure. Therefore, the boiling point of the liquid is much higher than 100°C. Autoclaves are used to sterilize surgical instruments, etc.

And vice versa, by reducing external pressure, we thereby lower the boiling point. By pumping air and water vapor out of the flask, you can make the water boil at room temperature ( Fig.11.3). As you climb mountains, the atmospheric pressure decreases, therefore 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 under these conditions.

Each liquid has its own boiling point, which depends on its saturated vapor pressure. 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 saturated vapor pressure of water is 101,325 Pa (760 mm Hg), and the pressure of mercury vapor is only 117 Pa (0.88 mm Hg). Mercury boils at a temperature of 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 important for boiling to increase the pressure of saturated vapor in the bubbles, and not to increase the pressure of the air in them?
3. How to make a liquid boil while cooling the vessel? (This question is not easy.)

G.Ya.Myakishev, B.B.Bukhovtsev, N.N.Sotsky, Physics 10th grade

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Why did people start boiling water before using it directly? That's right, to protect yourself from many pathogenic bacteria and viruses. This tradition came to the territory of medieval Russia even before Peter the Great, although it is believed that it was he who brought the first samovar to the country and introduced the ritual of leisurely evening tea drinking. In fact, our people used some kind of samovars back in ancient Rus' for preparing drinks from herbs, berries and roots. Boiling was required here mainly to extract useful plant extracts rather than for disinfection. After all, at that time it was not even known about the microcosm where these bacteria and viruses lived. However, thanks to boiling, our country was spared global pandemics of terrible diseases such as cholera or diphtheria.

Celsius

The great meteorologist, geologist and astronomer from Sweden originally used the value of 100 degrees to indicate the freezing point of water under normal conditions, and the boiling point of water was taken to be zero degrees. And after his death in 1744, no less famous person, botanist Carl Linnaeus and Celsius receiver Morten Stremer, inverted this scale to make it easier to use. However, according to other sources, Celsius himself did this shortly before his death. But in any case, the stability of the readings and understandable calibration influenced the widespread spread of its use among the most prestigious scientific professions at that time - chemists. And, despite the fact that, inverted, the scale mark of 100 degrees established the stable boiling point of water, and not the beginning of its freezing, the scale began to bear the name of its primary creator, Celsius.

Below the atmosphere

However, not everything is as simple as it seems at first glance. Looking at any phase diagram in P-T or P-S coordinates (entropy S is a direct function of temperature), we see how closely temperature and pressure are related. Likewise, water changes its values ​​depending on pressure. And any climber is well aware of this property. Anyone who has experienced altitudes above 2000-3000 meters above sea level at least once in their life knows how difficult it is to breathe at altitude. This is because the higher we rise, the thinner the air becomes. Atmospheric pressure drops below one atmosphere (below sea level, that is, below " normal conditions"). The boiling point of water also drops. Depending on the pressure at each height, it can boil at either eighty or sixty

Pressure cookers

However, it should be remembered that although most microbes die at temperatures above sixty degrees Celsius, many can survive at eighty degrees or more. That is why we achieve boiling water, that is, we bring its temperature to 100 ° C. However, there are interesting kitchen appliances that allow you to reduce time and heat liquid to high temperatures, without boiling it or losing mass through evaporation. Realizing that the boiling point of water can change depending on pressure, engineers from the USA, based on a French prototype, introduced the world to a pressure cooker in the 1920s. The principle of its operation is based on the fact that the lid is pressed tightly against the walls, without the possibility of steam escaping. Created inside high blood pressure, and water boils at more high temperatures. However, such devices are quite dangerous and often lead to explosions and serious burns to users.

Ideally

Let's look at how the process itself begins and goes through. Let us imagine an ideally smooth and infinitely large heating surface, where the heat distribution occurs evenly (the same amount of thermal energy is supplied to each square millimeter of the surface), and the surface roughness coefficient tends to zero. In this case, at n. u. boiling in a laminar boundary layer will begin simultaneously over the entire surface area and occur instantly, immediately evaporating the entire unit volume of liquid located on its surface. This ideal conditions, V real life This doesn't happen.

In reality

Let's find out what the initial boiling point of water is. Depending on the pressure, it also changes its values, but the main point here lies in this. Even if we take the smoothest pan, in our opinion, and bring it under a microscope, then in its eyepiece we will see uneven edges and sharp, frequent peaks protruding above the main surface. We will assume that heat is supplied uniformly to the surface of the pan, although in reality this is also not a completely true statement. Even when the pan is on the largest burner, the temperature gradient on the stove is distributed unevenly, and there are always local overheating zones responsible for the early boiling of water. How many degrees are there at the peaks of the surface and at its valleys? Peaks of the surface, with uninterrupted supply of heat, warm up faster than lowlands and so-called depressions. Moreover, surrounded on all sides by low-temperature water, they better transfer energy to water molecules. The thermal diffusivity coefficient of peaks is one and a half to two times higher than that of lowlands.

Temperatures

That is why the initial boiling point of water is about eighty degrees Celsius. At this value, the surface peaks supply enough of what is necessary for instantaneous boiling of the liquid and the formation of the first bubbles, visible to the eye, which timidly begin to rise to the surface. Many people ask what is the boiling point of water at normal pressure. The answer to this question can be easily found in the tables. At atmospheric pressure stable boiling is established at 99.9839 °C.

Boiling- this is vaporization that occurs simultaneously both from the surface and throughout the entire volume of the liquid. It consists in the fact that numerous bubbles float up and burst, causing a characteristic seething.

As experience shows, boiling of a liquid at a given external pressure begins at a well-defined temperature that does not change during the boiling process and can only occur when energy is supplied from the outside as a result of heat exchange (Fig. 1):

where L is the specific heat of vaporization at the boiling point.

Boiling mechanism: a liquid always contains a dissolved gas, the degree of dissolution of which decreases with increasing temperature. In addition, there is adsorbed gas on the walls of the vessel. When the liquid is heated from below (Fig. 2), gas begins to be released in the form of bubbles at the walls of the vessel. Liquid evaporates into these bubbles. Therefore, in addition to air, they contain saturated steam, the pressure of which quickly increases with increasing temperature, and the bubbles grow in volume, and consequently, the Archimedes forces acting on them increase. When the buoyant force becomes more power gravity of the bubble, it begins to float. But until the liquid is uniformly heated, as it ascends, the volume of the bubble decreases (saturated vapor pressure decreases with decreasing temperature) and, before reaching the free surface, the bubbles disappear (collapse) (Fig. 2, a), which is why we hear a characteristic noise before boiling. When the temperature of the liquid equalizes, the volume of the bubble will increase as it rises, since the saturated vapor pressure does not change, and the external pressure on the bubble, which is the sum of the hydrostatic pressure of the liquid above the bubble and the atmospheric pressure, decreases. The bubble reaches the free surface of the liquid, bursts, and saturated steam comes out (Fig. 2, b) - the liquid boils. The saturated vapor pressure in the bubbles is almost equal to the external pressure.

The temperature at which the saturated vapor pressure of a liquid is equal to the external pressure on its free surface is called boiling point liquids.

Since the saturated vapor pressure increases with increasing temperature, and during boiling it must be equal to the external pressure, then with increasing external pressure the boiling point increases.

The boiling point also depends on the presence of impurities, usually increasing with increasing concentration of impurities.

If you first free the liquid from the gas dissolved in it, then it can be overheated, i.e. heat above boiling point. This is an unstable state of liquid. Small shocks are enough and the liquid boils, and its temperature immediately drops to the boiling point.