Gas. Ideal gas

Liquefied natural gas or abbreviated LNG, as it is customary to call it in the energy industry (English resp. Liquefied Natural Gas, abbr. LNG) is ordinary natural gas cooled to -162°C (the so-called liquefaction temperature) for storage and transportation in liquid form. The liquefied gas is stored at the boiling point, which is maintained due to LNG evaporation. This method of storing LNG is due to the fact that for methane, the main component of LNG, the critical temperature is –83°C, which is much lower than the ambient temperature, and does not make it possible to store liquefied natural gas in high pressure tanks (for reference: the critical temperature for ethane is +32°C, for propane +97°C). For use, the LNG is evaporated to its original state without the presence of air. At ( return of the gas to its original vapor state) from one cubic meter of liquefied gas, about 600 cubic meters of ordinary natural gas are formed.

LPG temperature

The extremely low temperature of LNG makes it cryogenic liquid. As a general rule, substances with a temperature of -100°C (-48°F) or even lower are considered cryogenic and require special technologies for processing. For comparison, the lowest recorded temperature on Earth is -89.2°C (Antarctic), and in the settlement -77.8°C (Oymyakon village, Yakutia). The cryogenic temperature of liquefied natural gas means that contact with LNG can change the properties of contacting materials, which will subsequently become brittle and lose their strength and functionality. Therefore, special technologies are used in the LNG industry.

Chemical composition of LNG

Crude oil and natural gas are fossil fuels known as "hydrocarbons" because they contain chemical combinations of carbon and hydrogen atoms. The chemical composition of natural gas depends on where the gas is produced and processed. Liquefied natural gas represents mixture methane, ethane, propane and butane with a small amount of heavier hydrocarbons and some impurities, in particular nitrogen and sulfur complexes, water, carbon dioxide and hydrogen sulfide, which may exist in the feed gas, but must be removed before. Methane is the most important component, usually, although not always, more than 85% by volume.

LPG Density

Since LNG is a mixture, liquefied natural gas density varies slightly with its actual composition. Density of liquefied natural gas, is generally in the range of 430–470 kilograms per cubic meter, and its volume is approximately 1/600 of that of the gas at atmospheric conditions. This makes it about a third lighter than air. Another consequence of these facts is that LNG has a lower density than water, which allows it to float on the surface in the event of a spill and return to vapor rather quickly.

Other properties of LNG

Liquefied natural gas is odorless, colorless, non-corrosive, non-flammable and non-toxic. LNG is stored and transported at ultra-low temperatures at atmospheric pressure (no high pressures). When exposed to the environment, LNG evaporates quickly, leaving no traces on water or soil.

In his liquid form Liquefied natural gas does not have the ability to explode or ignite. At evaporation natural gas can ignite if it comes into contact with a combustion source and if the concentration of vapor in the air is between 5 and 15 percent. If the gas vapor concentration is less than 5 percent, then there is not enough vapor to start a fire, and if more than 15 percent, then there will be a lack of oxygen in the environment.

Benefits of LNG

1. The gas density increases hundreds of times, which increases the efficiency and convenience of storage, as well as transportation and energy consumption.
2. Liquefied natural gas - non-toxic cryogenic liquid, which is stored in a heat-insulated container at a temperature of –162°С. Large volumes of LNG can be stored at atmospheric pressure.
3. Possibility of intercontinental transportation of LNG by special vehicles, as well as transportation by rail and road transport in tanks.
4. Liquefied natural gas makes it possible to gasify facilities remote from main pipelines over long distances by creating an LNG reserve directly at the consumer, avoiding the construction of expensive pipeline systems.

From the consumer's point of view, the advantages of liquefied natural gas, based on it, also consist in the fact that LNG is not only a source of gas transported through gas pipelines, but also a source of NGL (broad fraction of light hydrocarbons- ethane, propane, butanes and pentanes), which are part of LNG and are released from LNG during regasification. These hydrocarbons are used as petrochemical raw materials and as a source of environmentally friendly fuel for various modes of transport (as well as in everyday life). In will be the selection of fraction With 2 + or With 3 +. Possibility to transport NGL as part of liquefied natural gas, it acts not only in favor of the consumer, but also solves the problems of the producer in terms of transportation NGL from the gas field.

Liquefied natural gas is a safe, environmentally friendly fuel with high energy performance and octane rating. LNG price at a cost to the consumer is lower than the price of liquefied petroleum gas, fuel oil, and even more so diesel fuel.

GAS. gaseous state called such a state of matter in which the forces acting between the molecules are extremely small and the dimensions of the molecules themselves are negligible compared with the gaps between them. Between collisions, gas molecules move in a straight line, evenly and completely randomly. When heated and rarefied, all gases tend to the limiting state of the so-called ideal, or perfect gas.

AT ideal gas the intermolecular forces are zero, and the volume of the molecules themselves is infinitesimal compared to the volume of the intermolecular space. The state of an ideal gas is that limiting diluted state of matter to which all bodies of nature aspire at sufficiently high temperatures and sufficiently low pressures; this is the special significance of the state of an ideal gas, which, moreover, is the most easily amenable to research and therefore the most fully studied. The substance that fills the interplanetary space in extreme rarefaction can be considered to be in the state of an ideal gas.

Gas pressure (p) is determined by the impact of gas molecules on the walls of the vessel. According to the kinetic theory, the average kinetic energy of gas molecules is proportional to the absolute temperature. In the kinetic theory, it is shown that an ideal gas strictly obeys the following equation of state, which relates three state parameters: v, T and p, of which two are independent, and the third is their function:

This equation ( Clapeyron's equation) contains in explicit form three basic laws of the state of an ideal gas:

1) Boyle-Mariotte law. At a constant temperature (T), the product (p ∙ v) for a given amount of ideal gas is a constant value (p ∙ v \u003d Const), i.e. the volume of an ideal gas (v) is inversely proportional to its pressure (p): isotherms of an ideal gas in coordinate system (v, p) are equilateral hyperbolas whose asymptotes are the coordinate axes.

2). At a constant (p), the volume of a given amount of ideal gas increases linearly with temperature:

(v 0 - volume at temperature \u003d 0 ° C, α - expansion coefficient of an ideal gas). Change (p) with temperature at v = Const obeys the same law:

(α) in equation (3) - pressure coefficient, numerically equal to the expansion coefficient (α) in equation (2) = 1/273.1 = 0.00367 - a value independent of the nature of the gas and the same for all ideal gases; p 0 - pressure at temperature \u003d 0 ° C. Introducing absolute temperature instead of temperature

we find instead of equations (2) and (3):

3) Avogadro's law. Equation (1) shows that gas constant R \u003d p 0 ∙ v 0 / 273.1 is proportional to the normal volume v 0 occupied by a given amount of gas under normal conditions (p 0 \u003d 1 Atm and t 0 \u003d 0 ° C \u003d 273.1 ° K), i.e. back is proportional to the gas density under normal conditions D 0 . According to Avogadro's law, with the same (p) and (T), all ideal gases contain in equal volumes (for example, equal to v 0) an equal number of molecules. Vice versa: an equal number of molecules (for example, 1 mol \u003d 1 gram molecule) of any gas in an ideal state occupies the same volume v 0 under normal conditions, regardless of the nature of the gas (1 mole of any substance contains N 0 = 6.06∙10 23 individual molecules - Avogadro's number). Found with great accuracy that normal molar volume any ideal gas (V 0) m is equal to 22.412 liters / mol. From here you can calculate the number of molecules in 1 cm 3 of any ideal gas under normal conditions: n0 \u003d 6.06 ∙ 10 23 / 10 3 ∙ 22.416 \u003d 2.705 ∙ 10 19 cm 3 (Loshmit number). Using equation (1), Avogadro's law is expressed in the fact that the gas constant R when calculated for 1 mole of any gas will be the same, regardless of the nature of the gas. That. R is a universal constant with dimension [ Job]/[weight][temperature] and expresses the work of expansion of 1 mole of an ideal gas when it is heated by 1 ° C at p \u003d Const:

this is the physical meaning of R.

find a numeric value

In other units, the R values ​​(per 1 mole) are:

In addition to the analyzed three laws, from the equation (1) of the state of an ideal gas in conjunction with the two laws of thermodynamics, the following basic laws also follow:

4) Joule's law. One of the general equations of thermodynamics

gives, together with equation (1), the following conditions for the internal energy U of an ideal gas:

i.e., U of an ideal gas is a function of only T (Joule's law); during isothermal expansion of an ideal gas, all absorbed heat is converted into external work, and during isothermal compression, all expended work is converted into heat released.

5) The heat capacities of an ideal gas at constant volume c v and at constant pressure c p are functions of T alone. Thermodynamics gives the general equations

but for an ideal gas (p) and (v) are linearly dependent on (T), according to the Gay-Lussac law (4) and (5); therefore, the right parts of equations (9) turn to 0 and

The heat capacities c p and c v are not independent of each other, but are related for an ideal gas by a simple condition:

arising from gas laws (R has the dimension of heat capacity), i.e., if c p and c v are related to 1 mole of an ideal gas, then they differ from each other by 2 (more precisely, by 1.986) - cal / mol ∙ deg.

In the kinetic theory, it is accepted, according to the principle of uniform distribution of energy, that for each degree of freedom of a gas molecule there is an energy k 0 ∙T / 2, and for 1 mole there is

(k 0 \u003d -R / N 0 is the gas constant calculated for 1 molecule - Boltzmann's constant). The number of degrees of freedom (i) is the number of types of mechanical energy that are independent of each other, which a gas molecule has. Then the energy of 1 mole

(approximately, assuming R = 2, c v = i, c p = i + 2).

In the theory of gas, the relation c p /c v = γ plays an important role; from equations (11) and (12):

In the simplest case monatomic gas(whose molecule consists of 1 atom, what are the noble gases and vapors of many metals) i is the smallest and equals 3: the entire energy of the molecule is reduced to the kinetic energy of its translational movements, which can be performed in three independent mutually perpendicular directions; then

and γ has the largest possible value: γ = 5/3 = 1.667. For diatomic gases(H 2 , O 2 , N 2 , CO and others) can be considered I \u003d 3 + 2 (two rotations around two mutually perpendicular axes perpendicular to the line connecting both atoms); then c v = 4.96 ≈ 5, cр = 6.95 ≈ 7 and γ = 7/5 = 1.40. For triatomic gas(Н 2 O, СO 2, H 2 S, N 2 O) i = 3+3 (rotation around three mutually perpendicular axes) and c v = 5.96 ≈ 6, cр = 7.95 ≈ 7 and γ = 4/ 3 = 1.33.

With further complication of the structure of the molecule, i.e., with an increase in i, c v and c p increase, and γ = 1 + 2/i and tends to 1. Table. 1 shows that everything said is in good agreement with the experimental data, that γ is always >1 and ≤1.667 and cannot be = 1.50 (for i = 4).

For monatomic gases, c v and c p, in accordance with the theory, practically do not change with temperature (for example, for Ar, the values ​​of c v and c p lie in the range from 2.98 to 3.00 between temperatures = 0 ° and 1000 ° C). Changes in c v and c p with temperature are explained in quantum theory. However, the heat capacities of gases that are close to ideal do not practically change over wide temperature ranges. Usually p and y are experimentally determined, and c v is calculated from these data.

real gases. All gases that actually exist are real gases b. or m. deviate from the laws of ideal gases, but the less, the higher the temperature and the lower the pressure. That. the laws of ideal gases are limiting for real gases. At ordinary temperatures, the deviation is least for gases whose critical temperatures are extremely low (the so-called constant gases: He, H 2 , N 2 , O 2 , air); for gases with a relatively high critical temperature and for vapors (a gas at a temperature below the critical temperature is called steam), the deviations are very significant. The reasons for deviations of real gases from gas laws are that: 1) intermolecular forces act in them; therefore, surface molecules are drawn into gases by forces, the resultant of which, calculated per unit surface and directed perpendicular to it, is called molecular (internal) pressure K; 2) not the entire volume of gas (v), but only part of it (v-b) gives freedom for the movements of molecules; part of the volume (b), covolum, as if occupied by the molecules themselves. If the gas were ideal, its pressure would be greater than the observed (p) by the value of K; therefore, the equation of state for a real gas will be written in the form

In this general equation, K and b may depend on T and v.

Van der Waals showed that in the simplest case, K \u003d a / v 2, and b is a constant value equal to four times the volume of the gas molecules themselves. Thus, the van der Waals equation has the form:

a and b, the van der Waals constants, as experience shows, still depend on T and v, and therefore equation (15) is only a first approximation; it reproduces well the qualitative shape of the isotherms of real gases.

In FIG. 1 are shown for CO 2 theoretical isotherms: S-shaped parts of these isotherms correspond to thermodynamically metastable states.

In FIG. Figure 2 shows the experimental isotherms for CO 2 : the S-shaped parts of the curves are replaced by straight parts; to the right of these parts, the curves correspond to gas (unsaturated vapor), to the left - to liquids, and the straight segments themselves - to the equilibrium of vapor and liquid. Equation (15), in full agreement with experience, shows that with increasing temperature, the dimensions of the straight segments on the isotherms become smaller and smaller (Fig. 2) and, finally, at a certain temperature equal to the critical temperature, the length of this segment becomes 0. At a temperature greater than At a critical temperature, a gas cannot turn into a liquid at any pressure: the liquid ceases to exist. That. the van der Waals equation covers two states - gaseous and liquid - and serves as the basis for the doctrine of the continuity of the transition between these two states. Critical temperatures for some gases have the following values: +360°C for H 2 O, +31°C for CO 2, -241°C for H 2 and -254°C for He.

Gas liquefaction. Any gas can be turned into a liquid at the proper pressure, having previously cooled it below the critical temperature. The pressures required for CO 2 liquefaction (in Atm) at different temperatures are given in Table. 2.

It is clear that these pressures are the pressures of saturated vapor of liquid carbon dioxide and the lower the lower the temperature.

In order to pre-cool the gas for liquefaction, in technical installations they use the Joule-Thomson effect, which consists in the fact that during adiabatic expansion (for example, with a sharp drop in pressure when the gas flows out of the hole), the internal energy of the gas increases by ΔU, and T changes by ΔT, and thermodynamically

In the case of ideal gases, ΔU = 0 and ΔT = 0 [because, according to equation (1), T∙dv/dT – v = 0].

For real gases, ΔТ ≠ 0, i.e. cooling or heating occurs, depending on whether T∙dv/dT – v ≠ 0 (Δp< 0). По уравнению Ван-дер-Ваальса,

(with sufficient approximation). That. at sufficiently high temperatures, all gases heat up during adiabatic expansion (ΔТ > 0, because a/R∙T< b), но с понижением температуры для каждого газа наступает inversion pointТ i determined by the condition

below which gases begin to cool during adiabatic expansion (a/R∙T> b at T< Т i). Для всех газов, кроме Н 2 и Не, Т i лежит выше обычных температур (так, для воздуха Т i соответствует +360°С), и потому газы могут быть сжижены по принципу Линде , без предварительного охлаждения. Для Н 2 инверсионная точка Т i - 80,5°С, а для Не - даже 15°К; поэтому Н 2 и Не для сжижения д. б. предварительно охлаждены ниже этих температур.

Relevant states. Critical temperature T to, pressure p to and volume v to m. b. expressed in terms of the van der Waals constants a, b and R as follows:

If we take critical values ​​for the units of measurement T, p and v, respectively, then instead of T, p and v, the state will be characterized by given values:

If we introduce θ, π, and ϕ into the van der Waals equation (15), then the constants a, b, and R cancel out, and we get reduced equation of state, with numerical coefficients

which does not contain quantities that depend on the nature of the substance. Equation (19) assumes, however, that the van der Waals equation is correct, and therefore deviations from it are often quite significant, especially in the case of associated substances. The doctrine of the corresponding states (the so-called states corresponding to the same θ, π and ϕ) makes it possible to find a large number of universal dependencies similar to equation (19).

Application of gases. Compressed and liquefied gases are used in technology wherever large quantities of gas are needed in a small volume; so, CO 2 is used for carbonation of waters, Cl 2 and phosgene - in the military chemical business, O 2 - for medical purposes, compressed air - for starting internal combustion engines. Liquefied gases (CO 2 and NH 3) are of particular importance in refrigeration, in refrigeration machines (for example, for producing artificial ice). Light gases (H 2, lighting gas, recently He) are used to fill balloons. Inert gases (N 2 and noble gases, especially Ar) are used to fill half-watt incandescent lamps. Of particular note is the use of gas for lighting or as a fuel: lighting, power, water gases and others.

Liquids can exist only at temperatures below the critical one. Therefore, to liquefy a gas, it must first be cooled below the critical temperature and then subjected to compression. As can be seen from Table XIII, gases such as oxygen, nitrogen, hydrogen, and especially helium require very low temperatures to liquefy.

Table XIII (see scan) Critical and boiling points (at atmospheric pressure) for some gases

One of the first industrial methods for liquefying gases (the Linde method, 1895) used the Joule-Thomson effect.

The scheme of the Linde machine is shown in Figure 6.21. Compressed by the compressor K and, as a result, somewhat heated, the gas passes through the cooler X, where it gives off heat to running water and cools to its original temperature. The gas then passes through the coil to a throttle valve (cock) and expands into receiver B with a pressure drop of about hundreds of atmospheres to one atmosphere. Immediately after the plant is started, the temperature drop is not sufficient to liquefy the gas. The slightly cooled gas is sent back to the compressor via a coil. Both coils are in close thermal contact (usually one coil is inserted into the other) in a counterflow heat exchanger. In the heat exchanger, the gas going to the compressor at a lower temperature cools the oncoming gas flow. Obviously, in the second cycle, the gas will approach valve A at a lower temperature than

this was during its first passage, and after throttling the temperature will drop even more. With each cycle, as a result of throttling and the action of the heat exchanger, the temperature of the gas will decrease more and more and eventually will drop so much that part of the gas, after expansion, turns into liquid and accumulates in receiver B, from where the liquid can be drained into the Dewar vessel through a valve

The described principle of countercurrent heat exchange is used in all machines for liquefying gases, although the design of such heat exchangers can be extremely diverse.

Another industrial method for liquefying gases (the Claude method, 1902) is based on the additional cooling of the gas when it does work. The compressed gas after the valve (Fig. 6.21) is sent to the piston machine (expander), where it, expanding, does the work of moving the piston due to the kinetic energy of the molecules (the expander is not shown in the figure). As a result, the effect of lowering the temperature of the gas becomes more significant than in the Linde machine. This method was improved by the Soviet scientist P. L. Kapitsa (1934), who instead of a piston expander used a small turbine (turbo expander) driven by a cooled gas (the expander rotor is small in size, and its weight is measured in only hundreds of grams).

At present, for the liquefaction of gases, in most cases, machines with expansion in expanders are used. When liquefying helium for pre-cooling in machines with turbo expanders, nitrogen is used instead of hydrogen, which significantly increases the productivity and economic efficiency of the device. In addition, with the same productivity, machines with turbo-expanders are several times smaller than machines operating according to the Linde scheme.

Instruction

Looks like liquefied natural gas(LNG) is a colorless and odorless liquid, 75-90% composed and possessing very important properties: in the liquid state, it is not combustible, nor is it aggressive, which is extremely important during transportation. The LNG liquefaction process has a character, where each new stage means compression by 5-12 times, followed by cooling and moving to the next stage. LNG becomes liquid upon completion of the last stage of compression.

If gas needs to be transported over very long distances, then it is much more profitable to use special vessels - gas carriers. From the place of gas to the nearest suitable place on the sea coast, a pipeline is being pulled, and a terminal is being built on the coast. There, the gas is highly compressed and cooled, turning it into a liquid state, and pumped into isothermal tanks of tankers (at temperatures of the order of -150 ° C).

This method of transportation has a number of advantages over pipeline transportation. Firstly, one of these in one flight can carry a huge amount of gas, because the density of a substance in a liquid state is much higher. Secondly, the main costs are not for transportation, but for loading and unloading the product. Thirdly, storage and transportation of liquefied gas is much safer than compressed gas. There can be no doubt that the share of natural gas transported in liquefied form will steadily increase compared to pipeline supplies.

Liquefied natural gas in demand in various fields of human activity - in industry, in road transport, in medicine, in agriculture, in science, etc. Liquefied gas We won due to the convenience of their use and transportation, as well as environmental friendliness and low cost.

Instruction

Before liquefying hydrocarbon gas and it must first be cleaned and removed water vapor. Carbonic gas removed using a three-stage molecular filter system. Purified in this way gas in small quantities it is used as a regeneration. Recoverable gas either incinerated or used to generate power in generators.

Drying occurs with the help of 3 molecular filters. One filter absorbs water vapor. Another dries gas, which goes further and passes through the third filter. To lower the temperature gas passed through a water cooler.

The nitrogen method involves the production of liquefied hydrocarbon gas and from any gas new sources. The advantages of this method include the simplicity of technology, the level of safety, flexibility, ease and low cost of operation. The limitations of this method are the need for a power source and high capital costs.

With a mixed method for the production of liquefied gas and a mixture of nitrogen and is used as a refrigerant. receive gas also from any source. This method features a flexible production cycle and low variable production costs. Compared to the nitrogen liquefaction process, capital costs are more significant here. A source of electricity is also needed.

Sources:

• What is gas liquefaction?
• Liquefied gas: receipt, storage and transportation
• what is liquefied gas

Natural gas is extracted from the bowels of the Earth. This mineral consists of a mixture of gaseous hydrocarbons, which is formed as a result of the decomposition of organic matter in sedimentary rocks of the earth's crust.

What are the ingredients in natural gas

80-98% natural gas consists of (CH4). It is the physicochemical properties of methane that determine the characteristics of natural gas. Along with methane, natural gas contains compounds of the same structural type - ethane (C2H6), propane (C3H8) and butane (C4H10). In some cases, in small quantities, from 0.5 to 1%, natural gas contains: (С5Н12), (С6Н14), heptane (С7Н16), (С8Н18) and nonane (С9Н20).

Natural gas also includes compounds of hydrogen sulfide (H2S), carbon dioxide (CO2), nitrogen (N2), helium (He), water vapor. The composition of natural gas depends on the characteristics of the fields where it is produced. Natural gas produced in pure gas fields consists mainly of methane.

Characteristics of natural gas constituents

All chemical compounds that make up natural gas have a number of properties that are useful in various industries and in everyday life.

Methane is a colorless, odorless, flammable gas that is lighter than air. It is used in industry and everyday life as a fuel. Ethane is a colorless, odorless, combustible gas that is slightly heavier than air. Basically, ethylene is obtained from. Propane is a poisonous, colorless and odorless gas. Butane is close to him in properties. Propane is used, for example, in welding work, in the processing of scrap metal. Liquefied and butane fill lighters and gas cylinders. Butane is used in refrigeration.

Pentane, hexane, heptane, octane and nonane -. Pentane is present in small amounts in motor fuels. Hexane is also used in the extraction of vegetable oils. Heptane, hexane, octane and nonane are good organic solvents.

Hydrogen sulfide is a poisonous colorless heavy gas, rotten eggs. This gas, even in small concentrations, causes paralysis of the olfactory nerve. But due to the fact that hydrogen sulfide has good antiseptic properties, it is used in small doses in medicine for hydrogen sulfide baths.

Carbon dioxide is a non-flammable, colorless, odorless gas with a sour taste. Carbon dioxide is used in the food industry: in the production of carbonated drinks to saturate them with carbon dioxide, to freeze food, to cool cargo during transportation, etc.

Nitrogen is a harmless colorless gas, odorless and tasteless. It is used in the production of mineral fertilizers, used in medicine, etc.

Helium is one of the lightest gases. It is colorless and odorless, non-flammable, non-toxic. Helium is used in various industries - for cooling nuclear reactors, filling stratospheric balloons.

The experimental fact of the cooling of a substance during evaporation has been known for a long time and has even been used in practice (for example, the use of porous vessels to preserve the freshness of water). But the first scientific study of this issue was undertaken by Gian Francesco Cigna and described in the work of 1760 "De frigore ex evaporation" ("On the cold due to evaporation").

Cigna proved that the faster the evaporation, the more intense the cooling, and Meran showed that if you blow on a wet bulb of a thermometer, the decrease in temperature will be greater than in the same experiment with a dry bulb of a thermometer. Antoine Beaumet (1728-1804) found that the evaporation of sulfuric ether cools more than the evaporation of water. Based on these facts, Tiberio Cavallo created a refrigerator in 1800, and Wollaston built his famous cryophore in 1810, which is still used today. Daniel's hygrometer was created on the basis of this device in 1820. The refrigeration machine became practically applicable only after 1859, that is, after Fernand Carré (1824-1894) published his method for obtaining ice by evaporating ether, which was later replaced by ammonia. In 1871, Carl Linde (1842-1934) described a refrigeration machine he had created in which cooling was achieved by gas expansion. In 1896, he combined this machine with the countercurrent heat exchanger described in physics courses, and this allowed him to obtain liquid hydrogen. The experimental results achieved by that time by physicists began to be introduced into industry.

The problem of gas liquefaction has a centuries-old history dating back to the second half of the 18th century. It all started with the liquefaction of ammonia by simple cooling, which was produced by van Marum, sulfuric anhydride by Monge and Clouet, chlorine by Northmore (1805) and the liquefaction of ammonia by the compression method proposed by Baccelli (1812).

Charles Cagnard de Latour (1777-1859) and Michael Faraday (1791-1867) simultaneously and independently made decisive contributions to the solution of this problem.

In a series of papers published in 1822 and 1823, Cañard de Latour described experiments carried out by him to determine the existence for a liquid (as it is felt intuitively) of a certain limiting expansion, beyond which, regardless of the applied pressure, all of it passes into a vapor state. To this end, de Latour placed a stone ball in a cauldron filled with alcohol one-third and began to gradually heat the cauldron. From the noise made by the ball turning inside the cauldron, de Latour came to the conclusion that at a certain temperature all the alcohol evaporated. The experiments were repeated with small tubes; air was removed from the tubes, and then they were filled to 2/5 with the investigated liquid (alcohol, ether, gasoline) and heated in a flame. As the temperature increased, the liquid became more and more mobile, and the interface between liquid and vapor became more and more indistinct, until at a certain temperature it completely disappeared and the entire liquid seemed to have turned into vapor. By connecting these tubes to a pressure gauge with compressed air, Cañard de Latour was able to measure the pressure established in the tube at the moment when the interface between liquid and vapor disappears, and the corresponding temperature. Contrary to popular belief, Cañard de Latour not only did not determine the critical temperature for water in these experiments, he did not even manage to completely evaporate the water, because the tubes always burst before the desired effect was achieved.

A more concrete result was contained in the experiments of Faraday, carried out in 1823 with bent glass tubes, the longer arm of which was sealed. In this arm, Faraday placed a substance that, when heated, was supposed to give the gas under study, then closed the second, short arm of the tube and immersed the tube in a cooling mixture. If, after doing this, the substance is heated in the long arm of the tube, then a gas is formed, the pressure of which gradually increases, and in many cases in the short tube Faraday liquefied the gas. Thus, by heating sodium bicarbonate, Faraday obtained liquid carbonic acid; in the same way, he received liquid hydrogen sulfide, hydrogen chloride, sulfuric anhydride, etc.

The experiments of de Latour and Faraday showed that a gas can be liquefied by subjecting it to high pressure. Many physicists began to work in this direction, in particular Johann Natterer (1821-1901). However, some gases (hydrogen, oxygen, nitrogen) could not be liquefied in this way. In 1850, Vertelo subjected oxygen to a pressure of 780 atm, but could not achieve liquefaction. This forced Vertelo to join the opinion of Faraday, who, confident that sooner or later it would be possible to obtain solid hydrogen, believed that pressure alone was not enough to liquefy certain gases, then called "permanent" or "indomitable".

In the same 1845, when Faraday expressed this consideration, Regnault, noticing that at low temperatures carbon dioxide has an anomalous compressibility, and when approaching 100 ° C begins to follow Boyle's law, put forward the assumption that for each gas there is a certain the temperature range where it obeys Boyle's law. In 1860, Regnault developed and modified this idea by Dmitry Ivanovich Mendeleev (1834-1907), according to which for all liquids there must be an “absolute boiling point”, above which it can exist only in a gaseous state, whatever the pressure.

The study of this question was resumed in 1863 in a new form by Thomas Andrews (1813-1885). In 1863, Andrews introduced carbon dioxide into a capillary tube, locking the volume of gas with a column of mercury. With the help of a screw, he arbitrarily set the pressure under which the gas was located, while gradually changing the temperature. Having achieved partial liquefaction of the gas by means of a mere increase in pressure and then slowly heating the tube, Andrews observed the same phenomena that Cañard de Latour had investigated 30 years before him. When the temperature of carbon dioxide reached 30.92°C, the interface between the liquid and the gas disappeared, and no amount of pressure could produce liquid carbon dioxide back. In his detailed work of 1869, Andrews proposed that the temperature of 30.92°C be called the "critical point" for carbon dioxide. By the same method, he determined the critical points for hydrogen chloride, ammonia, sulfuric ether, and nitric oxide. He proposed to retain the term "steam" for gaseous substances at a temperature below the critical point, and apply the term "gas" to substances at a temperature above the critical point. This point of view of Andrews was confirmed by the already mentioned experiments of Natterer, carried out by him from 1844 to 1855, in which permanent gases were subjected to pressure up to 2790 atm, without liquefying, and numerous similar experiments begun in 1870 by Emil Amaga (1841- 1915), in which pressures up to 3000 atm were achieved.

All these negative experimental results confirmed Andrews' hypothesis that permanent gases are substances for which the critical temperature is lower than the values ​​\u200b\u200breached at that moment, so that their liquefaction could be carried out using preliminary deep cooling, possibly with subsequent compression. This hypothesis was brilliantly confirmed in 1877 by Louis Calet (1832-1913) and Raoul Pictet (1846-1929), who, independently of each other, managed to liquefy oxygen, hydrogen, nitrogen, and air after a strong preliminary cooling. The works of Calhete and Pictet were continued by other physicists, but only the advent of the Linde refrigeration machine, which we have already mentioned, made liquefaction methods practically accessible, making it possible to obtain liquefied gases in large quantities and widely apply them in scientific research and in industry.

SPECIFIC HEAT CAPACITY OF GASES

Methods for determining the specific heat capacity were difficult to apply to gaseous substances due to the small specific gravity of gases and vapors. Therefore, at the beginning of the 19th century, the Paris Academy of Sciences announced a competition for the best method for measuring the specific heat of a gas. The prize was awarded to Francois Delaroche (? - 1813?) and Jacques Berard (1789-1869), who proposed to place a coil in the calorimeter, through which, at a known temperature, a gas would pass at a fixed pressure. This method was not actually new; it had been proposed 20 years earlier by Lavoisier. Be that as it may, the results obtained by Delaroche and Berard were presented in physics courses for half a century. The merit of these scientists, first of all, is that attention was drawn to the need to distinguish between specific heat capacities at constant pressure and at constant volume. The latter value is very difficult to measure because of the low heat capacity of the gas compared to the heat capacity of the reservoir containing it.

But a few years before the appearance of the works of Delaroche and Berard, research began on a curious phenomenon, noted by Erasmus Darwin (1731-1802) in 1788, and then in 1802 by Dalton, which consists in the fact that the compression of air causes it to heat up, and the expansion leads to to cooling. The beginning of the study of this phenomenon is usually considered the experience of Gay-Lussac (1807), repeated by Joule in 1845. Gay-Lussac connected two cylinders with a tube, just as Guericke did; one of the cylinders was filled with air, and the second was empty; from a filled cylinder, air could freely flow into an empty one. As a result, a decrease in the temperature of the first cylinder and an increase in the temperature of the second were found. This thermal behavior of air led us to believe that the specific heat at constant pressure must be greater than at constant volume, no matter what theory of the nature of heat we adhere to. Indeed, if the expanding gas cools, then by allowing it to expand during heating, it is necessary to impart additional heat to it in order to compensate for the cooling accompanying the expansion.

Based on these experimental facts, Laplace in 1816 came up with the brilliant idea that the well-known discrepancy between the value of the speed of sound, obtained from experience, and its theoretical value, obtained from Newton's law, can be explained by the change in temperature experienced by layers of air when alternating compressions and rarefactions. On the basis of these theoretical premises, Laplace corrected Newton's formula by introducing into it a coefficient equal to the ratio of specific heat capacities at constant pressure and at constant volume for air. Comparison of the experimental value of the speed of sound in air and the theoretical value obtained from Newton's formula made it possible to find the ratio of specific heat capacities. In this indirect way, physicists managed to obtain the first data on the value of this ratio and, thus, since the value of the specific heat at constant pressure was known, to estimate the specific heat of air at constant volume. A few years later (1819), Nicolas Clément (1779-1841) and Charles Desorme (1777-?) managed to directly determine the ratio of heat capacities , which, within experimental errors, coincided with that found by Laplace.

In 1829, as a result of subtle and painstaking research, Dulong determined the ratio of heat capacities for various gases, for which he caused sound in a tube using flows of various gases. These experiments led him to come to the conclusion that in gases and vapors under equal conditions (volume, pressure, temperature), the same amount of heat is formed with the same relative compression or expansion.

Note that Dulong's method was significantly improved in 1866 by Kundt (1839-1894), who introduced a special tube (this tube is now called Kundt's tube). The Kundt method is still considered one of the best methods for determining the ratio of specific heat capacities.

Compiled by Savelyeva F.N.