Characteristics of the main liquid on Earth: the physical and chemical properties of water. Structure and properties of water

The main substance that allows life to exist on the planet is water. It is essential in every situation. The study of the properties of liquids led to the formation of a whole science - hydrology. The subject matter of most scholars is physical and Chemical properties . They understand these properties as critical temperatures, crystal lattice, impurities and other individual characteristics of a chemical compound.

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Water Formula known to every student. These are three simple signs, but they are contained in 75% of the total mass of everything on the planet.

H2O- these are two atoms and one -. The structure of the molecule has an empirical form, so the properties of the liquid are so diverse, despite the simple composition. Each of the molecules is surrounded by neighbors. They are connected by one crystal lattice.

Simplicity of structure allows a liquid to exist in several states of aggregation. Not a single substance on the planet can boast of this. H2O is very mobile, it is second only to air in this property. Everyone is aware of the water cycle, that after it evaporates from the surface of the earth, it rains or snows somewhere far away. Climate regulated It is precisely due to the properties of a liquid that can give off heat, while itself practically does not change its temperature.

Physical properties

H2O and its properties depend on many key factors. The main ones are:

• Crystal cell. The structure of water, or rather its crystal lattice, is determined by the state of aggregation. It has a loose, but very strong structure. Snowflakes show a lattice in a solid state, but in the usual liquid state, water has no clarity in the structure of crystals, they are mobile and changeable.
• The structure of the molecule is a sphere. But the influence of gravity causes water to take the shape of the vessel in which it is located. In space, it will be geometrically correct.
• Water reacts with other substances, including those that have unshared electron pairs, among them alcohol and ammonia.
• Has high heat capacity and thermal conductivity heats up quickly and doesn't cool down for a long time.
• It has been known since school that the boiling point is 100 degrees Celsius. Crystals appear in the liquid when it drops to +4 degrees, but ice forms with an even greater decrease. The boiling point depends on the pressure in which H2O is placed. There is an experiment in which the temperature of a chemical compound reaches 300 degrees, while the liquid does not boil, but melts the lead.
• Another important property is surface tension. The water formula allows it to be very durable. Scientists have found that it will take a force with a mass of more than 100 tons to break it.

Interesting! H2O, purified from impurities (distilled), cannot conduct current. This property of hydrogen oxide appears only in the presence of salts dissolved in it.

Other features

Ice is unique condition which is characteristic of hydrogen oxide. It forms loose bonds that are easily deformed. In addition, the distance between particles increases significantly, making the density of ice much lower than that of liquid. This allows water bodies not to freeze completely in winter, keeping life under a layer of ice. Glaciers are a large reservoir of fresh water.

Interesting! H2O has a unique state called the triple point phenomenon. This is when it is in three of its states at once. This condition is possible only at a temperature of 0.01 degrees and a pressure of 610 Pa.

Chemical properties

Basic chemical properties:

• Divide water by hardness, from soft and medium to hard. This indicator depends on the content of magnesium and potassium salts in the solution. There are also those that are constantly in the liquid, and some can be disposed of by boiling.
• Oxidation and reduction. H2O affects the processes studied in chemistry that occur with other substances: it dissolves some, reacts with others. The outcome of any experiment depends on right choice the conditions under which it takes place.
• Influence on biochemical processes. Water main part of any cell, in it, as in an environment, all reactions in the body take place.
• In the liquid state, it absorbs gases that are inactive. Their molecules are located between the H2O molecules inside the cavities. This is how clathrates are formed.
• With the help of hydrogen oxide, new substances are formed that are not associated with the redox process. These are alkalis, acids and bases.
• Another characteristic of water is the ability to form crystalline hydrates. Hydrogen oxide remains unchanged. Among the usual hydrates, copper sulfate can be distinguished.
• If an electric current is passed through the connection, then molecules can be broken down into gases.

Importance for a person

A very long time ago, people understood the invaluable importance of liquid for all living things and the planet as a whole. . Without her man cannot live and weeks . What is the beneficial effect of this most common substance on Earth?

• The most important application is the presence in the body, in the cells, where all the most important reactions take place.
• The formation of hydrogen bonds favorably affects living beings, because when the temperature changes, the fluid in the body does not freeze.
• A person has long been using H2O for domestic needs, in addition to cooking, these are: washing, cleaning, bathing.
• No industrial plant can operate without fluid.
• H2O - source of life and health she is the cure.
• Plants use it at all stages of their development and life. With its help, they produce oxygen, a gas that is so necessary for the life of living beings.

In addition to the most obvious useful properties, there are still a lot of them.

The importance of water for humans

Critical temperature

H2O, like all substances, has a temperature that called critical. The critical temperature of water is determined by the method of its heating. Up to 374 degrees Celsius, the liquid is called vapor, it can still turn back into its usual liquid state, at a certain pressure. When the temperature is above this critical point, then the water as chemical element, turns into a gas irrevocably.

Application in chemistry

H2O is of great interest to chemists due to its main property - the ability to dissolve. Often, scientists purify substances with it, which creates favorable conditions for conducting experiments. In many cases, it is an environment in which pilot tests can be carried out. In addition, H2O itself participates in chemical processes, influencing one or another chemical experiment. It combines with non-metallic and metallic substances.

Three states

Water appears before people in three states, called aggregate. These are liquid, ice and gas. The substance is the same in composition, but different in properties. At

The ability to reincarnate is a very important characteristic of water for the entire planet, thus its circulation takes place.

Comparing all three states, a person more often sees a chemical compound in liquid form. Water has no taste and smell, and what is felt in it is due to the presence of impurities, substances dissolved in it.

The main properties of water in a liquid state are: a huge force that allows you to sharpen stones and destroy rocks, as well as the ability to take any form.

Small particles, when frozen, reduce the speed of their movement and increase the distance, therefore porous ice structure and less dense than liquid. Ice is used in refrigeration units, for various household and industrial purposes. In nature, ice brings only destruction, falling in the form of hail or avalanches.

Gas is another state that forms when the critical temperature of the water is not reached. Usually at temperatures over 100 degrees, or evaporating from the surface. In nature, these are clouds, fogs and vapors. Artificial gas formation played a large role in technological progress in the 19th century, when steam engines were invented.

The amount of matter in nature

75% - such a figure will seem huge, but this is all the water on the planet, even the one that is in different aggregate states, in living beings and organic compounds. If we take into account only liquid, that is, water located in the seas and oceans, as well as solid water - in glaciers, then the percentage becomes 70.8%.

Percentage distribution something like this:

• seas and oceans - 74.8%
• H2O fresh springs, distributed unevenly across the planet, in glaciers is 3.4%, and in lakes, swamps and rivers only 1.1%.
• On underground springs accounts for approximately 20.7% of the total.

Characteristics of heavy water

Natural substance - hydrogen occurs as three isotopes, in the same number of forms there is oxygen. This allows you to select, in addition to the usual drinking water also deuterium and tritium.

Deuterium has the most stable form, it is found in all natural sources, but in very small quantities. Liquid with such a formula has a number of differences from simple and light. So, the formation of crystals in it begins already at a temperature of 3.82 degrees. But the boiling point is slightly higher - 101.42 degrees Celsius. It has a higher density and the ability to dissolve substances is significantly reduced. In addition, it is denoted by another formula (D2O).

Living systems react on such a chemical compound is bad. Only some types of bacteria were able to adapt to life in it. The fish did not survive such an experiment at all. In the human body, deuterium can stay for several weeks, and then it is excreted without causing harm.

Important! Do not drink deuterium water!

The unique properties of water. - Just.

Conclusion

Heavy water has found wide application in the nuclear and nuclear industries, and ordinary water has found widespread use.

Content: It is necessary to distinguish, on the one hand, water and, on the other hand, the substances dissolved in it, which determine the chemical composition and mineralization of water. The geological fates of solvent and solute may follow their own separate paths. Water most often enters the earth's crust and from the atmosphere, and the solute is borrowed mainly from rocks and soils. Let us take water in its pure form, without salts, and consider those features of its structure and properties on which the dissolving power of water depends.

The composition of water. Water - a chemical compound of oxygen and hydrogen, which is usually denoted by the formula H 2 O. In fact, water has a more complex composition. The usual molecular weight of water is 18, but there are molecules with a molecular weight of 19, 20, 21, 22. These molecules consist of heavier hydrogen and oxygen atoms, having atomic weights of more than 1 and 16, respectively. Hydrogen has two stable isotopes: protium (H) and deuterium (D); ratio H: D =6800. In addition, tritium (T) is known - a radioactive isotope with a half-life of 12.5 years. Oxygen has three stable isotopes: O 16, O 17, O 18. Water molecules can consist of various stable isotopes H 2 O 16, HDO 16, D 2 O 16, H 2 O 18, HDO 18, D 2 O 18, H 2 O 17, HDO 18, D 2 O 17.

The isotopic variety of water in which protium is replaced by deuterium is called heavy water. However, neither light nor heavy water has been discovered in nature so far. Heavy water is currently prepared artificially in large quantities for various technical purposes. Heavy water differs from ordinary water not only in its physical properties, but also in its physiological effect on the body.

Deuterium (D) is of particular geochemical and practical interest. The electron shell of the deuterium atom, like protium, consists of one electron, but its nucleus - the deuteron - is about twice as heavy and consists of two particles - a proton and a neutron. Deuterium is used in modern nuclear technology as an explosive. In the future, it will be used as fuel in thermonuclear power plants. The reserves of thermonuclear energy of deuterium, available in the water of the earth's oceans, are approximately one hundred million superior to the reserves of energy of fossil fuels (coal, oil, gas, peat).

Natural waters of different genesis have different isotopic composition. One of the main reasons that create the differentiation of isotopes in natural waters is the process of evaporation The vapor pressure of heavy water is somewhat lower than the vapor pressure of ordinary water, and since the evaporation process is the main factor in the water cycle, the enrichment of water in heavy isotopes at evaporation sites and their depletion at condensation sites can cause a noticeable difference in water density.

The following pattern of distribution of hydrogen isotopes in surface and atmospheric waters has been established:

1. Fresh surface waters of rivers, lakes and other bodies of water, filled mainly due to precipitation, contain less deuterium than oceanic waters.

2 The isotopic composition of fresh surface waters is determined by the physical and geographical conditions of their location.

The structure of water. Back in the twenties of our century, on the basis of the doctrine of the polar structure of water molecules, the simplest ideas about the association of molecules in liquid water as a result of the interaction of dipoles were developed. These representations are as follows.

One of the features of the structure of the water molecule is the asymmetric arrangement of hydrogen atoms around the oxygen atom; they are located not in a straight line drawn through the center of the oxygen atom, but at a certain angle (Fig. 1). The centers of the nuclei of hydrogen atoms are located at a distance of 0.95 A from the center of the oxygen atom. The angle between the lines connecting the centers of oxygen and hydrogen atoms is 105 0 . The bond between oxygen and hydrogen atoms in a water molecule is carried out by electrons. Due to the asymmetry of the distribution of electric charges, the water molecule has polarity, i.e. has two poles - positive and negative, which, like a magnet, create force fields around it.

Thus, water molecules are characterized by dipole: structure (dipoles). They are depicted as ovals, the poles of which have electric charges opposite in sign. When approached sufficiently, water molecules begin to act on each other with their force fields . In this case, the positively charged pole of one molecule attracts the negatively charged pole of the other. As a result, aggregates of two, three, and apparently more molecules can be obtained (Fig. 2).

Such groupings of water molecules are called dihydrols (H 2 O) 2 and trihydrols (H 2 O). Consequently, single (monohydrols), double and triple molecules are simultaneously present in water . Their content varies with temperature. Ice is dominated by ternary molecules with the largest volume. As the temperature rises, the speed of the molecules increases, and the attractive forces between the molecules are insufficient to keep them close to each other. . In the liquid state, water is a mixture of dihydrols, trigdrols and monohydrols. As the temperature increases, the triple and double molecules break up, and at 100°C water consists mainly of monohydrols.

Chemically pure water has a number of properties that sharply distinguish it from other natural bodies.

1. When water is heated from 0 to 4 ° C, the volume of water does not increase, but decreases, and its maximum density is reached not at the freezing point (0 0 C), but at 4 0 C (more precisely, 3.98 0).

2. When water freezes, it expands, and does not shrink, like all other bodies, its density decreases.

3. The freezing point of water decreases with increasing pressure, and does not rise, as one would expect.

4. The specific heat capacity of water is extremely high compared to the heat capacity of other bodies.

5. Due to the high dielectric constant water has a greater dissolving and dissociating power than other liquids.

6. Water has the highest surface tension of all liquids - 75 erg / cm 2 (glycerin - 65, ammonia - 42, and all others are below 30 erg / cm 2), with the exception of mercury - 436 erg / cm 2.

Surface tension and density determine the height to which a liquid can rise in a capillary system when filtered through porous media.

The reason for these anomalous properties of water lies in the peculiarities of the structure of its molecules.

Water as a solvent. If you put water in the outer electric field, then the molecules of its iodine, under the action of the field, tend to settle down in space as shown in

This phenomenon is called orientational polarization, which substances with polar molecules have. The high polarity of water molecules is one of the most important reasons for its high activity in many chemical interactions. It also causes electrolytic dissociation in water, salts, acids and bases. It is also associated with the solubility of electrolytes in water.

Dissolution is not only a physical but also a chemical process. Solutions are formed by the interaction of solute particles with solvent particles. Water has the ability to dissolve many substances, that is, to give them homogeneous physico-chemical systems of variable composition (solutions). Dissolved in natural waters, salts are: mainly in a dissociated state, in the form of ions. In the solid crystalline state, ionic compounds consist of regularly arranged positive and negative ions. There are no molecules in this case.. Thus, for example, in halite, as determined by X-ray structural analysis, each Na + ion is surrounded by six C1 - ions, and each non C1 - is surrounded by six sodium ions. Ions interact with each other, they attract each other (ionic bond).

What is the dissolution mechanism? Water molecules, due to the peculiarities of their structure and the resulting force field around them, have the ability to attract molecules of other substances. . The process of dissolution consists precisely in the interaction of the particles of the dissolved substance with the particles of water. When a salt of some salt comes in contact with water, the nones that form its crystal lattice will be attracted by oppositely charged particles of water molecules. For example, when halite crystals are immersed in water, the sodium ion (cation) will be attracted by the negative pole, and the chlorine ion (anion) by the positive pole of the water molecule (Fig. 4 ). In order for the ions of the crystal lattice to break away from each other and go into solution, it is necessary to overcome the attractive force of this lattice. When salts are dissolved, such a force is the attraction of lattice ions by water molecules, characterized by the so-called hydration energy. If, in this case, the hydration energy is sufficiently large compared to the energy of the crystal lattice, the ions will be torn off from the latter and go into solution.

Depending on the nature of the substance, when it is dissolved, heat is usually released or absorbed. The ions of the dissolved substance attract and hold around themselves a certain number of water molecules, which form a shell called a gpdrate. Thus, in aqueous solution ions are hydrated, i.e. chemically bonded to water molecules

During the crystallization of many salts, part of the water of hydration is captured by crystal lattices. Gypsum CaSO 4 * 2H 2 O, mirabilite Na 2 SO 4 * 10H 2 O, bischofite MgCl 2 * 6H 2 O, astrakhanite Na 2 SO 4 * MgSO 4 * 4H 2 O, soda Na 2 CO 3 * 10H2O . Crystalline substances containing water molecules are called crystalline hydrates.

Strong electrolytes, when dissolved in water, completely dissociate into ions. These include almost all salts, many mineral acids, bases of alkali and alkaline earth metals. The dissociation of a strong electrolyte, such as NaCl, is represented by the equation

NaС1 Na + +С1 -

There are no NaCl molecules in a halite crystal. When dissolved, the crystal structure is destroyed, hydrated ions go into solution. There are no molecules in solution. Therefore, it is only conditionally possible to speak of non-dissociated molecules in solutions of strong electrolytes. It will rather be ion pairs (Na + +C1 -), i.e.

Oppositely charged ions located close to each other (closer to a distance equal to the sum of the radii of the ions). These are supposedly non-dissociated molecules, or, as they are called, quasi-molecules.

Weak electrolytes, when dissolved in water, only partially dissociate into ions. These include almost all organic acids, some mineral acids, such as H 2 CO, H 2 S, H 2 SіO 3, many metal bases. Water is a weak electrolyte.

In addition to electrolytes, there are also non-electrolytes in the solution, the molecules of which, although they have a hydrate shell, are "so strong that they do not decompose into ions (O 2, No. 2).

Depending on the size of the particles of the dissolved substance, true and colloidal solutions are distinguished. Solutions are called true when the solute is in the ionized state. According to the principle of electrical neutrality, an ionic solution always contains equal amounts of equivalents of cations and anions. Under natural conditions, ionic solutions are formed by dissolving simple salts.

Colloidal solutions are such solutions in which the substance is not in an ionized state, but in the form of groups of molecules, the so-called "colloidal particles". Particle sizes in colloidal solutions lie approximately in the range from 10 to 2000 A. In stable colloidal solutions, particles in most cases carry electric charges of different magnitudes, but the same sign for all particles of a given colloidal system. Colloidal solutions are called sols. Sols are capable of transforming into gels; turn into gelatinous masses as a result of enlargement of colloidal particles (coagulation process).

In nature, colloidal solutions can be organic or inorganic. The latter are formed mainly during the hydrolytic splitting of various silicates. During hydrolysis, silicates release the bases contained in them (alkali and alkaline earth metals), giving rise to true solutions. But, in addition, during hydrolysis, silicon, iron, aluminum and other metals pass into the solution, forming, for the most part, colloidal solutions.

Many substances enter into an exchange decomposition reaction with water, called hydrolysis. During hydrolysis, there is a shift in the equilibrium of water dissociation H O H + OH due to the binding of one of its ions by ions of the dissolved substance with the formation of a poorly dissociated or hardly soluble product. Therefore, hydrolysis is the chemical interaction of dissolved salt ions with water, accompanied by a change in the reaction of the medium. In view of the reversibility of hydrolysis, the equilibrium of this process depends on all those factors that generally affect the equilibrium of ion exchange. In particular, it strongly (sometimes almost completely) shifts towards the decomposition of salt, if the products of the latter (most often in the form of basic salts) are sparingly soluble.

Hydrolysis plays an important role in nature. For example, the main chemical form of mineral weathering in igneous rocks is hydrolysis.

Salt solubility. Solid, liquid and gaseous substances can dissolve in water. By solubility in water, all substances are divided into three groups: 1) highly soluble, 2) poorly soluble, and 3) practically insoluble. It must be emphasized that there are no absolutely insoluble substances.

The mineralization of natural waters is usually created by a few simple salts: chlorides, sulfides, bicarbonates of sodium, magnesium, calcium.

There are no NaCl molecules in the halite crystal. When dissolved, the crystal structure is destroyed, hydrated ions go into solution. There are no molecules in solution. Therefore, it is only conditionally possible to speak of non-dissociated molecules in solutions of strong electrolytes. These are the ion pairs (Na + Cl ), i.e. oppositely charged ions close to each other. These are non-dissociated molecules, but quasi-molecules.

Weak electrolytes, when dissolved in water, only partially dissociate into ions. These include almost all organic crystals, some mineral acids, such as H CO, H S, H SiO, many metal bases. Water is a weak electrolyte.

In addition to electrolytes, there are also non-electrolytes in the solution, the molecules of which, although they have a hydrated shell, are so strong that they do not decompose into ions (О, N).

Depending on the size of the particles of the solute, true and colloidal solutions are distinguished. Solutions are called true when the solute is in an ionized state.

The solubility of solids in water depends not only on their chemical nature, but also on temperature, pressure, and the presence of gases and impurities in it.

The solubility of sodium chloride changes little with increasing temperature from up to 60°C (the change in solubility is given in g per 100 mg of water). The solubility of sodium carbonate and sulfate increases greatly.

Temperature has a great influence on the solubility of silicic acid. In the silicic acid-water system, studied in the range from 0 to 200°, the dependence of solubility on temperature is linear. Under normal conditions, the solubility of silicic acid is very low.

Among the salts that decrease their solubility with increasing temperature is Ca SO 4 .

As is known, the solubility of a given salt decreases in the presence of another salt that has an ion of the same name with it, and, conversely, increases if non-similar ions are present in the solution. For example, the solubility limits of CaSO 4 in the presence of various salts vary greatly. If there is a large amount of sodium chloride in the solution (about 100 g / l), the solubility of CaSO 4 reaches 5-6 g / l

Of the main salts, the carbonates of alkaline earths have the lowest solubility, but it increases several times if the water contains carbon dioxide (CO 2) Dissolution proceeds according to the scheme:

CaCO 3 + H 2 O + CO 2 Ca (HCO 3) 2 Ca ++ +2HCO 3;

MgCO 3 + H 2 O + CO 2 Mg (HCO 3) 2 Mg ++ +2HCO 3.

These reactions are reversible and proceed until a certain equilibrium is reached. As a result of these reactions, calcium and magnesium bicarbonates appear in water. It should be noted that neither calcium bicarbonates nor magnesium bicarbonates exist in solid form. Mineralization of hydrocarbonate magnesium-calcium waters widely distributed in nature usually reaches 500-600 mg/l. In the presence of large amounts of CO 2, the solubility of Ca (HCO 3) 2 and Mg (HCO 3) 2 can exceed 1 g / l (carbonic mineral waters).

With increasing temperature, the solubility of calcium and magnesium bicarbonates decreases greatly and drops to 0 at 100 °. At high temperatures, these salts decompose with the release of CO 2 and the precipitation of carbonates

Ca (HCO 3) 2 → CaCO 3+H 2 O+CO 2;

Mg(HCO 3) 2 → MgCO 3+H 2 O+CO 2;

It follows from this that hydrocarbonate calcium and magnesium waters cannot exist in deep conditions, and, therefore, thermal waters of this composition do not exist.

The enrichment of waters with salts is accomplished not only by simple dissolution. Natural solutions are also formed during the hydrolytic splitting of certain minerals. Among the minerals that are insoluble directly in water, but capable of hydrolytically splitting, are various silicates-aluminosilicates, ferrosilicates, etc., which make up 75% of all minerals in the earth's crust. Under the influence of water and carbon dioxide during weathering, silicates give the base Na + , K + , Ca ++ , Mg ++ into the solution. These bases form, by combining with CO 2 , carbonic and bicarbonate salts or, under appropriate conditions, sulfate and chloride salts.

Main literature: OL 1 .

additional literature: DL 5.7.

Control questions:

1. What are the natural main isotopes?

2. What are the special qualities of water?

3. How does the process of halite dissolution take place?

4. How are substances divided and named by solubility?

The composition of water can be determined using the decomposition reaction electric shock. Two volumes of hydrogen are formed per one volume of oxygen (the volume of gas is proportional to the amount of substance):

2H 2 O \u003d 2H 2 + O 2

Water is made up of molecules. Each molecule contains two hydrogen atoms linked by covalent bonds to one oxygen atom. The angle between the bonds is about 105°:
O-H
H

Since oxygen is a more electronegative element (a strong oxidizing agent), the common electron pair of the covalent bond is shifted to the oxygen atom, a partial negative charge δ− is formed on it, and a partial positive δ+ is formed on hydrogen atoms. Neighboring molecules are attracted to each other by opposite charges - this causes a relatively high boiling point of water.

Water at room temperature- colorless transparent liquid. Melting point 0º C, boiling point at atmospheric pressure- 100° C. Pure water does not conduct electricity.

An interesting feature of water is that it has the highest density of 1 g / cm 3 at a temperature of about 4 ° C. As the temperature decreases further, the density of water decreases. Therefore, with the onset of winter, the upper freezing layers of water become lighter and do not sink down. Ice forms on the surface. The freezing of a reservoir to the bottom usually does not occur (besides, ice also has a density less than water and floats on the surface).

Chemical properties:

The main pollutants of natural water are sewage industrial enterprises containing compounds of mercury, arsenic and other toxic elements. Effluent from livestock complexes and cities may contain waste that causes the rapid development of bacteria. A great danger to natural water bodies is improper storage (which does not provide protection from atmospheric precipitation) or the use of fertilizers and pesticides washed into water bodies. Transport, especially water, pollutes water bodies with oil products and household waste thrown by unscrupulous people directly into the water.

To protect water, it is necessary to introduce closed water supply to industrial enterprises, complex processing of raw materials and waste, construction of treatment facilities, and environmental education of the population.

* Salt solutions are used for water electrolysis

2. Experience. Recognition of the salt of carbonic acid among the three proposed salts.

A qualitative reaction to carbonates is the interaction with acids, accompanied by a rapid release of carbon dioxide:

CaCO 3 + 2HCl \u003d CaCl 2 + H 2 O + CO 2

or, in ionic form:

CO 3 2− + 2H + = H 2 O + CO 2

It is possible to prove that it is carbon monoxide (IV) that is released by passing it through a solution of lime water, which causes it to become cloudy:

CO 2 + Ca (OH) 2 \u003d CaCO 3 ↓ + H 2 O

To recognize the salt of carbonic acid, add a little acid to all three test tubes (so that it does not overflow when “boiling up”). Where a colorless, odorless gas will be released, there is carbonate.

Let us first find out the structure of the thermodynamic precursor of water - ice. Thus, we will repeat the path of all water researchers. Each of them, trying to understand the structure of water, sooner or later came to the need to understand the structure of ice.

In 1910, the American physicist P. Bridgman and the German researcher G. Tamman discovered that ice can form several polymorphic crystalline modifications. Now 9 modifications of ice are known, they have different crystal lattices, different densities and melting points. The well-known ice is called "ice I", other modifications of ice exist at pressures exceeding 2000 atm. For example, ice III, which forms at a pressure of 2115 atm, is heavier than water, and ice VI (at a pressure of about 20,000 atm) melts at a temperature exceeding 80 °C. Under normal conditions, we can observe only ice I, and it has been studied most fully. Below we are talking exactly about him.

Each water molecule can form up to four hydrogen bonds if there are enough suitable neighbors nearby, and due to the property of cooperativity, each subsequent bond requires less energy to form, so it will be more likely to form than the previous one.

In ice, all molecules are linked by hydrogen bonds. In this case, four bonds of each molecule are locally organized into a tetrahedral structure, i.e. four nearby molecules are located at the vertices of a trihedral pyramid, in the center of which is the fifth water molecule.

Thus, the tetrahedral shape of an individual molecule is repeated in the crystal structure of ice. Perhaps a certain role here is played by the fact that the H-O-H angle of the H 2 O molecule is almost equal to the ideal tetrahedral angle of 109 °, and water molecules, as we know, are combined using hydrogen bonds, which they form precisely in O-H direction. These three-sided pyramids can also be combined into a kind of superstructure. In ice, such a complex three-dimensional superstructure of tetrahedra extends over the entire volume.

Starting from any oxygen atom, moving from neighbor to neighbor in hydrogen bonds, you can build an infinite number of different closed figures. All such figures are some kind of "corrugated" polygons, and the number of sides is always a multiple of six, and the shortest path from the molecule "to itself" passes along the sides of an ordinary hexagon. Therefore, the structure of ice is called hexagonal, or hexagonal.

If we forget about tetrahedra, we can see that the molecular structure of ice consists of zigzag layers, with each H 2 O molecule associated with three molecules of its own layer and one molecule of the neighboring layer. The number of neighbors of one molecule (in this case equal to four) is called the coordination number and is easily measured by the X-ray diffraction method. As you can see, the openwork network of hydrogen bonds turns the molecular structure of ice into a loose structure with a large number of voids.

If ice I is squeezed too much, it will change into other crystalline forms, and although its structure will change somewhat, the basic elements of the tetrahedral structure will remain. At moderate pressures (ice II, VI and IX), some of the hydrogen bonds break out of the tetrahedral structure (due to which the ice becomes somewhat denser), but any four nearest oxygen atoms are still combined by hydrogen bonds. Even at very high pressures (ice VIII and VII), the tetrahedral structure is preserved locally.

For the first time, the molecular structure of ice was established at the beginning of our century by the English scientist William Bragg, who developed the X-ray diffraction method for analyzing crystals. He found that each H 2 O molecule in ice was surrounded by four other molecules. But he was able to investigate precisely the molecular structure of ice, to establish how oxygen and hydrogen atoms are located in this structure, neither Bragg nor anyone else could at that time. Bragg used the X-ray diffraction method, which at that time allowed only relatively large atoms, such as oxygen or silicon, to be observed. Small atoms like hydrogen are not visible in x-ray diffraction analysis. Only at the end of the 40s of the 20th century, when new, more sensitive spectroscopic methods appeared, was it possible to establish the arrangement of hydrogen atoms in the structure of ice.

However, back in 1932, Bragg's student Professor Bernal was able to understand purely speculatively how oxygen and hydrogen atoms should be located in the molecular structure of ice.

Bernal proceeded from the configuration of the H 2 O molecule. He realized that it was the water molecule that determined the entire structure of ice. Bernal reasoned as follows: each hydrogen atom can "hook" only on one "foreign" oxygen atom, thereby linking two oxygen atoms ("own" and "foreign" atoms) with one hydrogen bond, therefore, each H 2 O molecule can connect using hydrogen bonds with four neighboring molecules, two of which form their own hydrogen atoms and two - atoms of neighboring molecules, and since the H 2 O molecule is "one-sided", this configuration should quickly fill the space, forming a tetrahedral structure.

These hypotheses were later confirmed by spectroscopic studies and are now known as the "Bernal-Fowler rules". Indeed, it turned out that each oxygen atom is associated with four hydrogen atoms located on the O-O line. It is connected with two "own" atoms covalent bond, and with two "strangers" - with the help of hydrogen bonding. Generally speaking, the terms "friend" and "alien" do not accurately describe the molecular life of ice. As it was found, not a single hydrogen is fixed in its place. Each hydrogen knows exactly only its O-O bond, but on this line it has two possible positions - near "its own" and near "foreign" oxygen atoms. In each of these positions, he spends on average half of his life time. If we denote, as is customary in chemistry, the valence bond with a dash, and the hydrogen bond with dots, then we can say that the reaction is continuously going on in ice:

O-H....O ↔ O....H-O

As you can see, the molecular life of ice is quite dynamic. But this applies only to hydrogen atoms, oxygen atoms sit firmly in their places and the distance in each couple O-O remains unchanged and equal to 2.76 A.

Obviously, the restlessness of hydrogen atoms must certainly affect the electrical and dielectric properties of ice. Ice has a fairly high electrical conductivity. Perhaps this feature of ice is explained by the fact that in the presence of an external electric field, jumps of hydrogen atoms become more directed.

Structure real ice is not absolutely ideal, in it, as in any other crystal, there are defects. The Danish researcher I. Bjerrum established that ice defects can be of two types: 1) there is not a single hydrogen atom on the O-O line (Bjerrum's A-defect); 2) there are two hydrogen atoms on the O-O line (D-defect). Of course, the energy of a defect is greater than the energy of a defect-free bond, so the defects do not sit on the same bond all the time, but migrate rather intensively throughout the entire ice structure. At the same time, they behave as if they are some particles of different signs. Two identical defects (for example, D-defects) will repel - after all, one defect leads to an increase in local energy, and even having two defects side by side is all the more energetically unfavorable. It is also intuitively clear that the two various defect will be attracted and at the meeting annihilate - destroy each other.

In ice, the concentration of defects is low - only one per 2.5 million molecules. So Bjerrum defects for ice are subtleties that are almost imperceptible for the structure of ice. It is a different matter in water, where the concentration of such defects increases by a factor of 25,000 and amounts to one defect per 100 molecules. This value is so significant that it becomes clear that Bjerrum defects play a significant role in water. An attempt was even made to describe water as ice with a high concentration of defects, which, in general, turned out to be not very consistent, but nevertheless, the theory constructed in this way was able to explain some phenomena.

Now let's move on to liquid water. The modern understanding of the molecular structure of water traces its history back to an article by British scientists Bernal and Fowler, which appeared in 1933 in the August issue of the newly created international journal on chemical physics, the Journal of Chemical Physics. This article remains one of the most remarkable milestones on the thorny path of understanding nature.

At that time, there was a rather simple - more philological than natural science - explanation of the anomalous properties of water. It was believed that water, the associated liquid, i.e. its molecules combine into large dehydrol supermolecules (H 2 O) 2 , (H 2 O) 3 , . . . (H 2 O) n, due to which water has anomalous properties. It was not at all clear why and how H 2 O molecules unite, how various associates are distributed over the volume of water. And most importantly, such an approach, generally speaking, did not explain the nature special properties water.

Trying to find his own understanding of the molecular structure of water, Bernal began by analyzing experimental facts. It cannot be said that at that time, in the 30s of the 20th century, these facts were enough, but still they were. Thanks to the brilliant research of William Bragg, the creator of X-ray diffraction analysis of crystals, the molecular structure of ice has been clarified. In addition to data on the structure of ice, Bernal had X-ray patterns of liquid water at his disposal, as well as the so-called radial distribution functions obtained using such X-ray patterns, i.e. the relative content of molecules located at certain distances from each other. In addition to purely experimental facts, Bernal had the opportunity, of course, to use ideas, hypotheses and assumptions, of which quite a lot had already accumulated by the beginning of the 1930s. However, the abundance of these ideas could hinder rather than help the development of the theory of water. Except, perhaps, for one old idea, dating back to the famous Wilhelm Roentgen, who suggested that the molecular structure of ice must somehow be repeated in the structure of liquid water. At one time, this idea was very popular among scientists, but all attempts to apply it to the description of the nature of the anomalous properties of water ended in failure. Even the simplest property of water - that it is heavier than ice - could not be explained with this idea. Moreover, it seemed that this feature of water simply contradicted it. Indeed, if we assume the existence of some strongly distorted structure of ice in the water, then the water should be lighter. Any violation of a clear structure, any disorder only increases the volume occupied by the structure. Therefore, such water should be lighter than ice.

In general, despite the beauty and temptation of the X-ray idea, no one was able to use it until the 1930s. It remained in the "bank of ideas" more as an aesthetic than a logical category, as a general statement that "water is a liquid that still retains the memory of the crystalline structure from which it originated" (formulation of the French physicist Clement Duval).

Analyzing the nature of water, Bernal spent a lot of time studying ice. He was already close to the theory of ice that we spoke about above. But the theory of ice itself, which is not capable of transforming into a theory of water, is of little value. But with water, everything was still unclear.

And then a chance intervened, which was pleased that in the rainy autumn of 1932 Professor Bernal went with a group of British scientists to Soviet Union. It also pleased chance that on the day of the departure of the British delegation to Moscow, a thick autumn fog descended. Aeroflot at that time did not pamper its customers with luxurious lounges, so Bernal had no choice but to wander around the airfield in the fog. Quite by chance, his companion on these walks turned out to be a very inquisitive person, Professor R. Fowler. “Most of all,” Bernal later recalled, “we were occupied with the fog that surrounded us, and it is natural that we were talking about it. The fog consists of water ... and Professor Fowler, a great connoisseur of thermodynamics, but not very knowledgeable in structural issues ", asked me to explain the structure of water, as I understand these problems. And then I thought about it again - in the light of our Moscow discussions." The walk of the two professors lasted more than twelve hours and turned out to be very fruitful, they managed to find a simple and beautiful solution to the problem of water. A few months later, the joint work of Bernal and Fowler appeared in print and became the basis of modern understanding of the molecular nature of water.

In telling Fowler about water, Professor Bernal mentioned and old idea X-ray, in which few people believed. Quite unexpectedly, they found an extremely important argument in favor of this idea. It was obtained by the "from simple" method. “What would happen to water,” Fowler asked, if it did not have a molecular structure? For example, what would be the density of such water? In such water, each H 2 O molecule must be surrounded by at least six neighbors, as in any close packing. It can be calculated that the density of such water would not be 1 g/cm 3 , but 1.8 g/cm 3 . Since at no temperature does the density of real water come close to this figure, it follows that in liquid water at any temperature there is some kind of molecular structure, most likely similar to the molecular structure of ice. It is this structure that keeps water molecules from being tightly packed.

Later, this assumption was confirmed by X-ray diffraction analysis, with the help of which it was possible to establish that the so-called "coordination number" of water (ie, the average number of neighbors of any molecule) is 4.4. Since the coordination number of ice is 4, the number of neighbors of the "statistically average" H 2 O molecule during the transition from the solid to the liquid state increases by only 0.4 neighbors. Therefore, out of every 10 water molecules, 8 are still surrounded by four neighbors, and two new molecules will appear near the other two.

Yes, but what about the anomalous behavior of ice during melting? After all, above we seem to have come to the conclusion that the distortion of the structure should lead to a decrease in the density of any substance. Discussing this contradiction, Bernal and Fowler finally came to the conclusion that when ice melts, it is not a distortion that occurs, but a rearrangement of the structure, while the long-range order of ice is destroyed, but inside small regions the molecular crystal-like structure is preserved. At that time, it was already known that such a rearrangement could lead to an increase in density. Bernal and Fowler in their paper referred to X-ray diffraction data for tridymite and quartz, which are very close to the corresponding data for ice and water. Tridymite and quartz are two different crystalline states of silica SiO 2 . The chemical composition of quartz and tridymite is the same, the molecular structures are also the same - both in quartz and in tridymite the molecules form tetrahedral structures. But the density of quartz is approximately 10% greater than the density of the tridymnt. Why is it the same structure, the same molecules, but the density is different? Bernal and Fowler knew the answer to this question. Since both oxygen and silicon are fairly large atoms, they are clearly visible on X-ray patterns, so all the subtleties of the structures of these crystals were already clarified in the 1930s. These subtleties consist in the fact that the distances between the nearest molecules in these crystals are the same, but the distance to the next (not nearest) neighbors is different for them, i.e. the first coordination spheres are the same for them, and the size of the second sphere for quartz is 4.2 A, and for tridymite - 4.5 A. This explains the differences in the density of quartz and tridymite.

If we recall that, firstly, ice also has a tetrahedral structure and, secondly, that the density of ice and water differ by 9%, then it is easy to understand the confidence of Bernal and Fowler that the structure of ice is similar to the structure of tridymite, and the structure water is similar to the structure of quartz. Not all the details of their theory have stood the test of time, more sophisticated theories later appeared, but their article in the Journal of Chemical Physics remains one of the most important milestones in the theoretical path of knowledge of water.

As is often the case, the Bernal-Fowler theory turned out to be correct only in its methodological part, and many of its details were not confirmed by further experiments. In particular, no quartz-like structures have been found in liquid water. But the idea of ​​water as a liquid with a highly developed openwork frame found more and more evidence.

The indisputable achievement of the 20th century was the clear understanding that the structure of ice is somehow preserved in water, or, using the formulation of Clement Duval, water remembers its origin. But why does she remember, while other liquids lack this ability? After all, ice (if you forget that it does not exist in "its own" temperature range), in general, is a fairly ordinary crystal. The fact that it has a special molecular structure is not so strange. All crystals form some (sometimes surprising) structures. But when melted, they give rise to quite trivial, ordinary liquids. Ice also melts and also creates liquid, but it is unusual. Why? To answer this question, let us recall that the molecules of most substances are held in the nodes of their crystal structures by rather weak van der Waals or electric forces. Molecules of H 2 O are held in the hexagonal structure of ice by hydrogen bonds, which differ significantly from van der Waals and electrostatic interactions. Hydrogen bonds are much stronger and, most importantly, their action is strictly directed in space. The last property leads to the fact that the hydrogen bond is only destroyed "immediately" when ice melts, it cannot gradually "deteriorate" before finally breaking. This is a very important difference between ice and other crystals. After all, when a crystal is heated, first of all, the thermal movement of individual molecules increases, which gradually deviate farther and farther from the entire node of the ideal crystal structure. And this is where the effect of directionality of hydrogen bonds comes into play. Let us assume that all crystal molecules sit in the nodes of an ideal structure. And suddenly one molecule jumps out of its node and moves away from it for some distance. In an ordinary substance, this molecule still retains a connection with its neighbors in the crystal lattice. Of course, the adhesion between them worsens, the energy of interaction increases, but the connection remains. If such an event occurs in ice, then the restless molecule will necessarily break all its hydrogen bonds, it cannot "slightly" deviate from the crystal lattice site, while retaining all its hydrogen bonds. After all, the hydrogen bonds of its neighbors are stretched to a very definite point in space, and if the molecule leaves this point, then thereby it loses the ability to "close" its two protons and two unshared electrons. At first glance, it may seem that just water should quickly forget its crystalline past. It turns out that H 2 O molecules "break" with their past immediately and irrevocably. Strictly speaking, this should be the case if a large number of molecules in ice could break all their hydrogen bonds at once. But in order for such an event to occur in the molecular life of ice, it is necessary to concentrate in one place at once a rather large (on a molecular scale) energy.

A single water molecule cannot gradually accumulate energy in order to break away from its neighbors upon reaching a certain energy level. To use the well-known physical lexicon, one can say that each ice molecule sits in a deep energy well with completely sheer edges. It is very difficult to jump out of such a hole, and if a molecule that has jumped out “stumbles”, it will immediately find itself at the bottom, in the structure perfect ice. Therefore, firstly, the probability of breaking hydrogen bonds is small, and secondly, having released only one H 2 O molecule from the crystal structure, ice immediately pays a rather large energy tribute to the kinetic melting processes and thus can retain a significant number of molecules in the crystal structure.

Energy pits, in which other substances are located, have a different appearance. Between the states corresponding to a crystal and a liquid, there is whole line intermediate states. Therefore, the molecules of ordinary substances can gradually accumulate energy, passing from one intermediate well to another. If any molecule loses some of its energy, then it will not be at the very bottom of the well, but may linger in some intermediate state. As a result, all crystal molecules are involved in the melting process rather quickly. The average energy of the molecules gradually increases, while the individual fluctuations in the energy of the molecules are not too large. If we depict the melting of an ordinary crystal in a certain phase-energy space, then it will be possible to see that during melting all the molecules are kept in a rather compact group. In fact, each point of such a space denotes the energy level of molecules. At the beginning of melting, all points will merge into one solid point corresponding to the crystalline state. In the process of melting ordinary matter, this point will creep up, gradually blurring and disintegrating into separate points. Then the central point will break up into smaller points, which, in turn, will also break up, and this process will end with the formation of a large, relatively dense swarm of points with a center corresponding to the liquid state. The picture of ice melting in this interpretation will look completely different. The peculiarity of the energy profile of ice molecules makes it possible to a large number H 2 O molecules during melting maintain a crystalline hexagonal structure of hydrogen bonds, only a small number of water molecules are actually involved in the melting process at any given time. At the beginning of melting, all molecules "sit" at the energy level corresponding to the state of ice. As the ice heats up, individual molecules break out of the crystal structure and immediately find themselves at the energy level of molecules without hydrogen bonds. There is a continuous exchange between these two levels; As the ice heats up, the number of molecules leaving the ice structure increases, while the number of those returning decreases. But even after the complete completion of melting, a fairly large part of the hydrogen bonds that existed in ice are retained in water as well.

The picture of ice melting described above is an idealization corresponding to the so-called two-structure model of water, i.e. a model in which only two states of H 2 O molecules are allowed - either completely free monomers or completely included in the hexagonal structure. In this regard, the question may arise: is such a mixture of monomers and a hexagonal lattice admissible? Recall that the structure of ice is loose, there are many voids in it, the atoms are located quite spaciously. Each cavity is surrounded by six H 2 O molecules, and each molecule is surrounded by six cavities that form continuous microscopic channels. The author of one of the first physical theories of water, the Soviet scientist O. Samoilov, calculated the size of the cavities and found that one water molecule could easily fit in them without touching or destroying the main framework of hydrogen bonds. Back in the 1940s, Samoilov suggested that during the melting of ice, part of the hydrogen bonds break, free H 2 O monomers appear, which partially fill the cavities of the hydrogen frame.

In 1952, the American scientists Heggs, Hasted and Buchanan managed to establish, using data on the dependence of the dielectric properties of water on temperature, that at 25 ° C in liquid water, 67% of all H 2 O molecules retain all four hydrogen bonds, 23.2% retain three hydrogen bonds. bonds, 7.6% - two hydrogen bonds, and only 0.2% - completely free molecules. Undoubtedly, the real structure of water is more complicated than that assumed by two-structure models, however, due to their simplicity, they are quite clear and are suitable as a "zero" approximation.

Other theories of the molecular state of water have also been proposed. For example, the English physicist D.Zh. Popl assumed that during the melting of ice, hydrogen bonds do not break at all, but somehow "bend". Professor Bernal, developing his idea, built a new theory of water, according to which H 2 O molecules form small closed rings of four, five or more molecules. But the vast majority of these rings, Bernal believed, consist of only five molecules, since the H-O-H angle in a water molecule is close to 108 ° - the angle of a regular pentagon.

L. Pauling in 1952 suggested that the structure of water is similar to the structure of clathrate hydrates of the Cl 2 10H 2 O type. Eyring put forward the theory of meaningful structures, which suggests that there are two crystal-like structures in water: ice I and ice III. Hydrogen bonds in the structure Ice III somewhat compressed and slightly curved, so this ice is 20% denser than ice I.

G. Nameti and X. Sheraga suggested that each water molecule can be in one of five permissible energy states, determined by how many hydrogen bonds it forms (0, 1, 2, 3 or 4). It is assumed that the molecules are collected in ice-like "swarms". Having done the usual analysis for statistical mechanics, Nameti and Sheraga found the number of water molecules in separate swarms that form 4, 3 and 2 hydrogen bonds. The molar volume of the system obtained in this way has a minimum at 4 °C, other parameters are also in good agreement with the experimental results. However, the theory of Namet and Sherag, as well as the two-structure model, contradicts a whole series of spectroscopic data. This is a common shortcoming of all theories that assume the existence of clearly distinct structures in water. IN real water, there appears to be a wide and continuous range of different molecular structures.

All theories (here we have mentioned only a few) are more or less consistent with the observed experimental data, but for each of them, sooner or later, facts were discovered that they could not explain. This, of course, does not mean that the theories are wrong. Each of them represented a certain degree of approximation to the true real picture of the physical state of water and worked for the future final theory.

With the advent of computers and the ability to simulate a variety of processes on them, it was possible to drastically reduce the number of reliable theories. With the help of such experiments, it was possible to determine exactly what proportion of water molecules retains all four hydrogen bonds, which - three, two, one, and how many completely free monomer molecules are in water. The figure shows a histogram of the distribution of hydrogen bonds in water at 10 °C obtained using a computer experiment.

As you can see, in water there is a fairly significant part of all types of molecules - from completely free to midnight bound. The histograms for other temperatures are similar, but at higher temperatures the maximum of the histogram (which in the case of 10 °C is at 2.3 hydrogen bonds per molecule) shifts to lower values ​​of the number of hydrogen bonds.

It turned out that both pentagons and hexagons are formed with equal success in water, without any preference for one over the other. This, by the way, means that hydrogen bonds can stretch and bend. The result obtained in this way crossed out all the models of "icebergs", which postulated that water is a sea of ​​completely free molecules, in which more or less large fragments of ice structures float. Although clusters with 1, 2, 3 ... number of hydrogen bonds are present, their share is small. Since the ice structures form only hexagons, such a campaign, of course, completely excludes the possibility of the appearance of pentagonal structures in the water.

Summarizing the results of numerous computer experiments, we can say that the topology of the molecular structure of water cannot be interpreted in the form of any hexagonal structure of ice with randomly broken hydrogen bonds. Moreover, this structure is a single entity in any volume of water. Computer experiments have shown that the network of hydrogen bonds is above the "critical percolation threshold". This means that in any volume of water there will always be at least one continuous chain of hydrogen bonds that permeates the entire volume of water.

How now, in the light of the results of computer experiments, can one imagine the physical nature of water? At the molecular level, water appears to be a randomly organized three-dimensional network of hydrogen bonds. Locally, this network tends to a tetrahedral configuration. This means that the nearest neighbors of the average water molecule are mainly located at the vertices of the tetrahedral pyramid surrounding the water molecule. The network contains a significant number of highly strained hydrogen bonds, and it is these bonds that play a fundamental role in the emergence of special anomalous properties of water. Any water molecule whose bonds are sufficiently strained can quickly change its entire immediate environment by switching its strained bonds to new neighbors. All this leads to the fact that the general topology of the entire network of hydrogen bonds in water is extremely variable and diverse. As ice melts, the distinct but loose tetrahedral structure is replaced by a less defined but more compact network of hydrogen bonds. The increase in density occurs due to the formation of more compact local structures (for example, the transition to pentagons from hydrogen bonds) and due to the bending of hydrogen bonds. When melt water is heated, the transition to more compact structures dominates up to 4 °C, after which processes associated with conventional thermal expansion prevail.

Fluid structure

The use of the term "structure" to describe ice is understandable, ice is a crystal and, of course, has an internal structure. But what is the structure of a liquid? "Isn't the lack of structure - fluidity - the defining quality of a liquid?" Bernal wrote. It turns out that the liquid has a structure, and not one, but several. It's all about the time scale.

If a coordinate system is associated with any fixed water molecule, then for an observer located in this system, the structure of water will depend on the characteristic time scale with which he will observe the molecular life of water. Water has two characteristic time parameters. Like any substance, be it a liquid or a solid, there is a period of oscillation of an individual molecule τ υ. For water, this value is 10 -13 s. In a liquid, in addition to the period of oscillations of molecules around their equilibrium position τ υ , there is one more characteristic time - the time of "settled life" τ D , i.e. the average time of existence of a given local environment of one molecule. For water τ D ~ 10 -11 s, i.e. before jumping to a new place, the water molecule makes 100 vibrations in one place.

These two parameters break the timeline into three regions, each of which has its own fluid structure. If the observer uses a sufficiently small time scale, i.e. will look for a time much less than τ υ, then he will see chaotically scattered molecules, among which it is difficult to discern any order. However, this random arrangement of molecules is called instantaneous, or M-structure.

To understand why this disorder is still called a structure, the observer needs to move to a longer time scale. But not too much, more precisely, more than τ υ , but less than τ D . At this time interval, real molecules will no longer be visible, the observer will be able to see only the points around which they perform their oscillations. It turns out that these points in the water are located quite regularly and form a clear structure, called the K-structure, which means "vibrationally averaged."

The M- and K-structures of water are similar to those of ice. To see the differences between these structures in water and ice, you need to observe them a little longer, i.e. with a characteristic time much greater than τ D . The pattern observed in this case is called the D-structure - diffusion-averaged. Unlike ice, the D-structure of water is completely blurred due to frequent jumps of water molecules over long distances (these jumps constitute the process of self-diffusion of water molecules). The D-structure is formed by diffusion averaging of K-structures and cannot be described by any special arrangement of points in space. An outside observer sees that, in fact, no D-structure of a liquid exists (note that it is the D-structure, as a complete statistical averaging of an ensemble of molecules, that determines the thermodynamic properties of water.).

Nevertheless, the D-structure exists and can be seen. An observer who is on a certain water molecule will see that his own molecule moves randomly throughout the entire volume of water, each time finding itself in a more or less ordered environment. He will see that most often "his" molecule will be surrounded by four other molecules of H 2 O, sometimes there will be five neighbors, sometimes six, on average, as we know, there will be 4.4 of them. Thus, the D-structure of water can be considered a picture seen by an observer.

This approach to describing the structure of water is most often used in the interpretation of spectroscopic data, because various spectroscopic methods - X-ray, NMR, dielectric relaxation, and neutron Raman scattering - are able to "read" molecular data with different characteristic resolution times.

The movement of molecules is usually proved by Brownian motion. A drop of water, in which very light particles of a solid insoluble substance float, is examined under a microscope and the particles are observed to move randomly in the body of water. Each such particle consists of many molecules and does not exhibit spontaneous motion. The particles experience shocks from the moving water molecules, which cause them to change direction all the time, which means that the water molecules themselves move randomly.

The importance of water for plant life

Lecture 10. Water exchange.

1. The importance of water for plant life

2. Structure and properties of water

3. Water exchange in a plant cell

3.1. Forms of water in plant cells

3.2. water potential. Osmosis. Transport of water in a plant cell

4. Osmotic absorption of water

5. Mechanisms for moving water

6. Top and bottom end motors

7. The movement of water through the vessels

8. Effect of water deficit on physiological processes

9. Features of water exchange of different ecological groups of plants

In plant tissues, water makes up 70-95% of the building mass. The role of water in the whole organism is diverse. Consider the functions of water in biological objects:

Water environment unites all parts of the body into a single whole. In the body of a plant, water is a continuous medium throughout, from the water taken up by the roots to the leaves evaporating water into the atmosphere.

Water is the most important solvent and medium for biochemical reactions;

Water is involved in the ordering of structures in cells, it is part of protein molecules, determining their conformation;

Water is a metabolite and a direct participant in biochemical reactions. For example, during photosynthesis, water is an electron donor, it is necessary for hydrolysis, for the synthesis of substances.

Water is the main component in the transport system of plants;

Water is a thermoregulatory factor, it protects plants from sudden temperature fluctuations;

Water is a shock absorber under mechanical influences;

Thanks to the phenomena of osmosis and turgor, it ensures the elastic state of cells (all plants, according to their ability to regulate the volume of moisture contained in them, are divided into poikilohydrothermal and homeohydrothermal. Poikilohydrothermal - cannot regulate the volume of water in the body, for example, algae, aquatic plants, etc. Homeohydrothermal plants can regulate the amount of water in the body through stomata).

Water can be in three aggregate states: solid, liquid and gaseous. In each of these states, the structure of water is not the same. During instant freezing with the help of liquid nitrogen, the water molecules do not have time to form a crystal lattice and the water acquires a solid glassy state (vitrification state). This property of water allows you to freeze living organisms without damage. The crystalline state of water is characterized by a wide variety of forms (for example, snowflakes).

2.1. Physical properties of water.

1. Density.

At 4 about C and a pressure of 1 atm. one cm3 of water weighs one gram. Those. the density of water is 1. When freezing, the volume of water increases by 11%.

2. Boiling and freezing points.

At a pressure of 1 atm. the boiling point of water is 100 o C, the freezing point is 0 o C. With increasing pressure, the freezing point decreases every 130 atm. 1 o C, and the boiling point increases.

3. Melting heat

The heat of melting ice is 0.335 kJ/h. Ice at normal pressure can have a temperature from -1 to -7 o C. The heat of vaporization of water is 2.3 kJ / h.

4. Heat capacity.

The heat capacity of water is 5-30 times higher than that of other substances. Heat capacity - the amount of heat required to raise the temperature by 1 o C. This feature of water is explained by the adhesion of molecules to each other (cohesion) due to hydrogen bonds.

5. Surface tension and adhesion.

On the surface of water (due to the ability of molecules to cohesion) surface tension is created. Water also has the property of adhesion (sticking), which is necessary when water rises against gravitational forces.