Pure water is a colorless transparent liquid. The density of water during its transition from a solid to a liquid state does not decrease, as in almost all other substances, but increases. When water is heated from 0 to 4°C, its density also increases. At 4°C, water has a maximum density, and only with further heating does its density decrease.

If, with a decrease in temperature and during the transition from a liquid to a solid state, the density of water changed in the same way as it happens with the vast majority of substances, then when winter approaches, the surface layers natural waters cooled down. to 0°C and sink to the bottom, making room for warmer layers, and this would continue until the entire mass of the reservoir would have acquired a temperature of 0°C. Further, the water would begin to freeze, the resulting ice floes would sink to the bottom and the reservoir would freeze to its entire depth. At the same time, many forms of life in water would be impossible. But since water reaches its greatest density at 4 °C, the movement of its layers, caused by cooling, ends when this temperature is reached. With a further decrease in temperature, the cooled layer, which has a lower density, remains on the surface, freezes, and thereby protects the underlying layers from further cooling and freezing.

Of great importance in the life of nature is the fact that water. has an abnormally high heat capacity, therefore, at night, as well as during the transition from summer to winter, the water cools slowly, and during the day or during the transition from winter to summer it also heats up slowly, thus being the temperature regulator on the globe.

Due to the fact that when ice melts, the volume occupied by water decreases, pressure lowers the melting point of ice. This follows from Le Chatelier's principle. Indeed, let. ice and liquid water are in equilibrium at 0°C. With increasing pressure, the equilibrium, according to Le Chatelier's principle, will shift towards the formation of that phase, which at the same temperature occupies a smaller volume. This phase is in this case liquid. Thus, an increase in pressure at 0 ° C causes the transformation of ice into a liquid, which means that the melting point of ice decreases.

The water molecule has an angular structure; the nuclei included in its composition form an isosceles triangle, at the base of which there are two protons, and at the top - the nucleus of the oxygen atom, Internuclear O-N distances close to 0.1 nm, the distance between the nuclei of hydrogen atoms is approximately 0.15 nm. Of the eight electrons that make up the outer electron layer of the oxygen atom in the water molecule, two electron pairs form covalent O-N connections, and the remaining four electrons are two unshared electron pairs.

The oxygen atom in the water molecule is in the -aea?eaecaoee state. Therefore, the HOH bond angle (104.3°) is close to tetrahedral (109.5°). The electrons that form O-H bonds are shifted to the more electronegative oxygen atom. As a result, hydrogen atoms acquire effective positive charges, so that two positive poles are created on these atoms. Centers negative charges lone electron pairs of the oxygen atom, which are in hybrid orbitals, are displaced relative to the nucleus of the atom and create two negative poles.

The molecular weight of vaporous water is 18 and corresponds to its simplest formula. However, the molecular weight of liquid water, determined by studying its solutions in other solvents, turns out to be higher. This indicates that in liquid water there is an association of molecules, i.e., their combination into more complex aggregates. This conclusion is also confirmed by the anomalously high values ​​of the melting and boiling points of water. The association of water molecules is caused by the formation of hydrogen bonds between them.

In solid water (ice), the oxygen atom of each molecule is involved in the formation of two hydrogen bonds with neighboring water molecules according to the scheme in which hydrogen bonds are shown by a dotted line. The diagram of the volumetric structure of ice is shown in the figure. The formation of hydrogen bonds leads to such an arrangement of water molecules, in which they are in contact with each other with their opposite poles. The molecules form layers, each of which is associated with three molecules belonging to the same layer, and with one from the adjacent layer. The structure of ice belongs to the least dense structures, there are voids in it, the sizes of the least dense structures, there are voids in it, the dimensions of which somewhat exceed the dimensions of the molecule.

When ice melts, its structure is destroyed. But even in liquid water, hydrogen bonds between molecules are preserved: associates are formed - like fragments of the ice structure - consisting of a larger or smaller number of water molecules. However, in contrast to ice, each associate exists for a very short time: the destruction of some and the formation of other aggregates is constantly taking place. In the voids of such "ice" aggregates, single water molecules can be placed; in this case, the packing of water molecules becomes denser. That is why when ice melts, the volume occupied by water decreases, and its density increases.

As the water heats up, the fragments of the ice structure in it become less and less, which leads to a further increase in the density of water. In the temperature range from 0 to 4°C, this effect prevails over thermal expansion, so that the density of water continues to increase. However, when heated above 4°C, the effect of increased thermal motion of molecules prevails and the density of water decreases. Therefore, at 4°C, water has a maximum density.

When water is heated, part of the heat is spent on breaking hydrogen bonds (the energy of breaking a hydrogen bond in water is approximately 25 kJ/mol). This explains the high heat capacity of water.

Hydrogen bonds between water molecules are completely broken only when water passes into steam.

Aggregate states of water

The physical properties of water are anomalous, which is explained by the above data on the interaction between water molecules. Water is the only substance on Earth that exists in nature in all three states of aggregation - liquid, solid and gaseous.

Density of water in solid and liquid state

The melting of ice at atmospheric pressure is accompanied by a decrease in volume by 9%. The density of liquid water at a temperature close to zero is greater than that of ice. At 0°C, 1 gram of ice occupies a volume of 1.0905 cubic centimeters, and 1 gram of liquid water occupies a volume of 1.0001 cubic centimeters. And the ice floats, which is why water bodies usually do not freeze through, but are only covered with an ice cover.

The temperature coefficient of volumetric expansion of ice and liquid water is negative at temperatures below - 2100 C and + 3.980 C, respectively.

Heat capacity of water

The heat capacity during melting almost doubles and in the range from 00 C to 1000 C is almost independent of temperature

Melting and boiling points of water in comparison with other hydrogen compounds of elements of the main subgroup YI group of the periodic table

Water has abnormally high melting and boiling points in comparison with other hydrogen compounds of elements of the main subgroup of group VI of the periodic table.

Water Status Diagram

The state diagram (or phase diagram) is graphic image dependences between the quantities characterizing the state of the system and phase transformations in the system (transition from a solid state to a liquid state, from a liquid to a gaseous state, etc.). State diagrams are widely used in chemistry. For one-component systems, state diagrams are usually used showing the dependence phase transformations on temperature and pressure; they are called P-T state diagrams.

The figure shows in a schematic form (without strict adherence to scale) a diagram of the state of water. Any point on the diagram corresponds to certain values ​​​​of temperature and pressure.

The diagram shows those states of water that are thermodynamically stable at certain temperatures and pressures. It consists of three curves that delimit all possible temperatures and pressures into three regions corresponding to ice, liquid and vapor.

Let's consider each of the curves in more detail. Let's start with the OA curve (Fig. 3), which separates the vapor region from the liquid state region. Imagine a cylinder from which air is removed, after which a certain amount of pure, free from dissolved substances, including gases, water is introduced into it; the cylinder is equipped with a piston, which is fixed in a certain position. After some time, some of the water will evaporate and saturated steam will be above its surface. You can measure its pressure and make sure that it does not change over time and does not depend on the position of the piston. If you increase the temperature of the entire system and again measure the pressure saturated steam, it turns out that it has increased. By repeating these measurements at various temperatures, we find the dependence of the pressure of saturated water vapor on temperature. The OA curve is a graph of this dependence: the curve points show those pairs of temperature and pressure values ​​at which liquid water and water vapor are in equilibrium with each other - coexist. The OA curve is called the liquid-vapor equilibrium curve or the boiling curve. The table shows the values ​​​​of saturated water vapor pressure at several temperatures.

fig.3(top)

Let us try to realize in the cylinder a pressure different from the equilibrium one, for example, less than the equilibrium one. To do this, release the piston and raise it. At the first moment, the pressure in the cylinder will indeed drop, but soon the equilibrium will be restored: an additional amount of water will evaporate and the pressure will again reach the equilibrium value. Only when all the water has evaporated can a pressure less than equilibrium be realized. From this it follows that the points lying on the phase diagram below or to the right of the OA curve correspond to the vapor region. If you try to create a pressure that exceeds the equilibrium, then this can only be achieved by lowering the piston to the surface of the water. In other words, the points of the diagram lying above or to the left of the OA curve correspond to the region of the liquid state.

How long do the regions of the liquid and vapor state extend to the left? Let's outline one point in both areas and we will move from them horizontally to the left. This movement of the points on the diagram corresponds to the cooling of a liquid or vapor at constant pressure. It is known that if you cool water at normal atmospheric pressure, then when it reaches 0 ° C, the water will begin to freeze. Carrying out similar experiments at other pressures, we arrive at the OS curve separating the region of liquid water from the region of ice. This curve is the equilibrium curve solid state- liquid, or melting curve, - shows those pairs of temperature and pressure values ​​at which ice and liquid water are in equilibrium.

Moving horizontally to the left in the vapor area (in the lower part of the diagram), we will similarly arrive at the 0V curve. This is the solid-state-vapour equilibrium curve, or the sublimation curve. It corresponds to those pairs of temperature and pressure values ​​at which ice and water vapor are in equilibrium.

All three curves intersect at point O. The coordinates of this point are the only pair of temperature and pressure values. at which all three phases can be in equilibrium: ice, liquid water and steam. It is called the triple point.

The melting curve has been studied up to very high pressures. Several modifications of ice have been found in this region (not shown in the diagram).

On the right, the boiling curve ends at critical point. At the temperature corresponding to this point, the critical temperature, the quantities characterizing physical properties liquid and vapor become the same, so that the distinction between liquid and vapor disappears.

The existence of a critical temperature was established in 1860 by D. I. Mendeleev, studying the properties of liquids. He showed that at temperatures above the critical one, a substance cannot be in a liquid state. In 1869, Andrews, studying the properties of gases, came to a similar conclusion.

Critical temperature and pressure for various substances different. So, for hydrogen = -239.9 °N, = 1.30 MPa, for chlorine = 144 ° C, = 7.71 MPa, for water = 374.2 ° C, = 22.12 MPa.

One of the features of water that distinguishes it from other substances is the decrease in the melting point of ice with increasing pressure. This circumstance is reflected in the diagram. The OC melting curve on the state diagram of water goes up to the left, while for almost all other substances it goes up to the right.

The transformations that occur with water at atmospheric pressure are reflected in the diagram by points or segments located on the horizontal corresponding to 101.3 kPa (760 mm Hg). So, the melting of ice or the crystallization of water corresponds to point D, the boiling of water corresponds to point E, the heating or cooling of water corresponds to segment DE, etc.

State diagrams have been studied for a number of substances of scientific or practical importance. In principle, they are similar to the considered diagram of the state of water. However, the state diagrams of various substances may have features. Thus, substances are known triple point which lies at a pressure greater than atmospheric pressure. In this case, heating the crystals at atmospheric pressure does not lead to the melting of this substance, but to its sublimation - the transformation of the solid phase directly into a gaseous one.

Heavy water

During the electrolysis of ordinary water, which contains, along with HO molecules, also an insignificant amount of DO molecules formed by the heavy isotope of hydrogen, HO molecules are predominantly decomposed. Therefore, during long-term electrolysis of water, the residue is gradually enriched in DO molecules. From such a residue after repeated repetition of electrolysis in 1933 For the first time, it was possible to isolate a small amount of water consisting of almost 100% D O molecules and called heavy water.

According to its properties, heavy water differs markedly from ordinary water (table). Reactions with heavy water proceed more slowly than with ordinary water. Heavy water is used as a neutron moderator in nuclear reactors.

Isotopic composition

There are nine stable isotopic varieties of water. Their average content in fresh water is as follows:

1H216O - 99.73%, 1H218O - 0.2%,

1H217O - 0.04%, 1H2H16O - 0.03%. The remaining five isotopic species are present in water in negligible amounts.

“There is nothing softer and weaker than water, but still nothing better for working hard and strong things.”

This paradox was formulated by the Chinese sage Lao Tzu in the ancient text Tao-Te-King, or Scripture on Morality. Indeed, water's ability to wash, soothe and nourish contrasts with unstoppable power, exemplified by Niagara Falls, the Grand Canyon (carved over the centuries by the Colorado River) and the tsunami.

Equally paradoxically, water is both utterly familiar—it makes up about two-thirds of our body and covers three-quarters of the planet—and utterly mysterious. Although you think you know her very well, many water properties you will be very surprised. And some of them are so strange that they are still not fully understood by science.

Downhill race

A logical person will assume that hot water will take longer to cool to 0 degrees Celsius and freeze than cold water. But the strange thing is that this is not always true. In 1963, a Tanzanian student named Erasto Mpemba noticed that in reality hot water freezes faster than cold water when two bodies of water are exposed to the same sub-zero conditions.

And no one knows why.

The only assumption is that the Mpemba effect results from a heat circulation process called convection. In the container, warm water rises, displacing cold water and creating an "insulated top". The scientists speculate that convection may somehow speed up the cooling process, allowing warmer water to freeze faster than cold water, no matter how much mercury it has to travel before freezing.

slippery substance

a century and a half scientific research never gave an answer why you can fall on the ice. Scientists agree that a thin layer of liquid water on top solid ice becomes the cause of slipperiness, and the mobility of the fluid makes it difficult to move, even if the years are thin. But there is no consensus as to why ice, unlike most solids has this layer.

Theorists suggest that it is the process of sliding, that is, contact with ice, that causes its surface to melt. Others believe that the liquid layer exists even before the sliding object appears, and that it is formed due to the internal movement of surface molecules.

Undoubtedly, you are looking for the culprit, lying on your back and seething with anger, but, unfortunately, he has yet to be found.

Aquanaut

On Earth, boiling water creates thousands of tiny bubbles of steam. In space, one giant oscillating bubble is created.

Fluid dynamics are so complex that physicists could not imagine what would happen to boiling water in zero gravity until an experiment was conducted on board in 1992. spaceship. After that, physicists decided that the simplified form of boiling in space is obviously associated with the absence of convection and lifting force Both of these phenomena are gravity. On Earth, these effects cause the bubbling that we see in a teapot.

floating liquid

When a drop of water hits a surface much hotter than its boiling point, it can float on the surface for much longer than you would expect. This is the Leidenfrost effect, and it comes from the fact that when the bottom layer of the droplet evaporates, gaseous molecules the water in this layer has nowhere to go, and their presence isolates the remainder of the drop and prevents it from touching the hot surface. Thus, the drop exists for several seconds before completely evaporating.

Extraordinary Shell

At times, water seems to defy the laws of physics, holding itself back from disintegrating, even though gravity or even the pressure of heavy objects is trying to tear it apart.

It is surface tension, the property that makes the outer layer of a body of water (and some other liquids) behave like a flexible shell. Surface tension occurs due to the fact that water molecules are weakly bonded to each other. Due to this, the surface molecules experience internal effort from the molecules below them. Water will remain intact until the force that tears it apart overcomes the force of these weak bonds and breaks through the surface.

For example, in the photo above, a paper clip rests on the surface of the water. Although metal is denser than water and therefore must sink, surface tension prevents the paperclip from breaking through the surface of the water.

boiling snow

When there is a huge difference in temperature between water and outside air, an amazing effect occurs - say, if you pour a pot of boiling water (100 degrees Celsius) into air at minus 34 degrees Celsius, then the boiling water will instantly turn into snow and scatter.

Explanation: Extremely cold air is very dense, the distance between its molecules is so small that there is not enough space left for the transport of water vapor. Boiling water, on the one hand, emits steam very actively. When it is thrown into the air, it breaks up into droplets, which creates even more space for the vapor to spread. This presents a problem. emitted more steam than can hold air, and so it spreads out, attaching itself to microscopic particles in the air, like soda or calcium, and forming crystals. This is how snowflakes are formed.

Empty space

Although the solid state of almost any substance is denser than the liquid state, since the atoms in solids are usually tightly packed together, this does not hold true for H2O. When water freezes, its volume increases by almost 8 percent. This is a strange property that allows ice cubes and even huge icebergs to float.

When water cools to the freezing point, there is less energy to cause the molecules to stick together, and therefore they can form stronger hydrogen bonds with their neighbors and gradually become fixed. This same process causes all liquids to solidify. And, as in other solids, the bonds between ice molecules are indeed shorter and stronger than in liquid water; the difference is that the hexagonal structure ice crystals leaves a lot of empty space, making ice generally less dense than water.

Excess volume can sometimes be seen in the form of ridges on top of ice cubes in your freezer. These protrusions are made up of excess water squeezed out of the cube by freezing (and expanding) ice. In a container, water freezes from the sides and bottom to the center and top, and the ice expands towards the center.

One of a kind

As the saying goes, no two snowflakes are the same. Indeed, throughout the history of snow exploration, every beautiful structure has been absolutely unique. And here's why: a snowflake is born in the form of a simple hexagonal prism. As she falls, she encounters non-repeating conditions that change their shape, including different temperatures, humidity levels and Atmosphere pressure. These variable factors are sufficient to ensure that crystal formation never occurs twice in the same pattern.

And what is most interesting about snowflakes is that all six of their branches grow in perfect synchrony, creating hexagonal symmetry, because each branch experiences the same conditions as all the others.

Where is she from?

The exact origin of water on our planet, which covers about 70 percent of the surface, is still a mystery to scientists. They suspect that any water that accumulated on the planet's surface during its formation over 4.5 billion years would have evaporated due to the intense heat of the young Sun. This means that the water we have now must have come later.

How? During a period called the Late Heavy Bombardment about 4 billion years ago, massive objects, possibly from other systems, fell to Earth and the planets. solar system. It is possible that such objects were filled with water, and these collisions could deliver huge volumes of this substance to our planet.

Comets — clumps of ice and rock with tails of evaporating ice, orbiting the Sun in long orbits — could very well be the remnants of what fell to the planet. However, there is a problem: remote studies of the water evaporating from several large comets have revealed that they are composed of water of a different type of H2O (containing a heavier isotope of hydrogen) than Earth's, therefore such comets cannot be the source of all our wonderful water.

The anomalous physical properties of water are so everyday and natural that we usually do not even suspect of their existence, completely forgetting that these properties are a gift from nature to all life on Earth.

Much has been written about water. Write scientists of different specialties - physicists, chemists, geologists, biologists, astronomers. There is even a certain tradition in writing statues about water to begin the story with a description of the unusual, anomalous properties of this liquid.

The melting and boiling point of water

The most surprising and blissful property of water for living nature is its ability to be a liquid under "normal" conditions. Molecules of compounds very similar to water (for example, H 2 S or H 2 Se molecules) are much heavier, but form a gas under the same conditions. Thus, water seems to contradict the laws of the periodic table, which, as you know, predicts when, where and what properties of substances will be close.

In our case, it follows from the table that the properties hydrogen compounds elements (called hydrides) located in the same vertical columns should change monotonically with increasing mass of atoms. Oxygen is an element of the sixth group of this table. In the same group are sulfur S (with an atomic weight of 32), selenium Se (with an atomic weight of 79), tellurium Te (with an atomic weight of 128) and pollonium Po (with an atomic weight of 209). Consequently, the properties of the hydrides of these elements should change monotonously when passing from heavy elements to lighter ones, i.e. in the sequence H 2 Po → H 2 Te → H 2 Se → H 2 S → H 2 O. Which is what happens, but only with the first four hydrides. For example, the boiling and melting points rise as the atomic weight of the elements increases. In the figure, the crosses mark the boiling points of these hydrides, and the circles mark the melting points.

As the atomic weight decreases, the temperatures decrease in a perfectly linear fashion. The area of ​​existence of the liquid phase of hydrides becomes more and more "cold", and if the oxygen hydride H 2 O were a normal compound, similar to its neighbors in the sixth group, then liquid water would exist in the range from -80 ° C to -95 ° C. At higher temperatures, H 2 O would always be a gas. Fortunately for us and all life on Earth, water is anomalous, it does not recognize a periodic pattern, but follows its own laws.

This is explained quite simply - most of the water molecules are connected by hydrogen bonds. It is these bonds that distinguish water from the liquid hydrides H 2 S, H 2 Se, and H 2 Te. If they were not, then the water would boil already at minus 95 ° C. The energy of hydrogen bonds is quite high, and they can be broken only at a much higher temperature. Even in gaseous state big number H 2 O molecules retain their hydrogen bonds, uniting into dimers (H 2 O) 2 . Fully hydrogen bonds disappear only at a water vapor temperature of 600 °C.

Recall that boiling consists in the fact that vapor bubbles form inside a boiling liquid. At normal pressure, pure water boils at 100 "C. If heat is supplied through the free surface, the process of surface evaporation will be accelerated, but volumetric evaporation characteristic of boiling does not occur. Boiling can also be carried out by lowering the external pressure, since in this case the pressure vapor equal to the external pressure is achieved at a lower temperature.At the top is very high mountain the pressure and, accordingly, the boiling point are so low that the water becomes unsuitable for cooking food - the required water temperature is not reached. With a high enough pressure, water can be heated enough to melt lead (327°C) and still not boil.

In addition to the super-high boiling points of melting (and the latter process requires too much heat of fusion for such a simple liquid), the very range of existence of water is anomalous - one hundred degrees, by which these temperatures differ - a rather large range for such a low molecular weight liquid as water. The limits of allowable values ​​of hypothermia and overheating of water are unusually large - with careful heating or cooling, water remains liquid from -40 ° C to +200 ° C. This extends the temperature range in which water can remain liquid to 240 °C.

When ice is heated, its temperature first rises, but from the moment a mixture of water and ice is formed, the temperature will remain unchanged until all the ice has melted. This is explained by the fact that the heat supplied to the melting ice is primarily spent only on the destruction of crystals. The temperature of melting ice remains unchanged until all crystals are destroyed (see latent heat of fusion).

Density of water and ice

Vital for the entire biosphere is the ability of water to decrease rather than increase its density when it freezes (as happens with almost all other substances). Bismuth behaves like water in this respect, but it is one of the extremely rare exceptions to general rule. G. Galileo was the first to pay attention to this unusual property of water. Indeed, during the transition of a liquid to a solid state, the molecules of a substance should seem to be located closer, and the substance itself should become denser. Normally things behave like this. But water is an exception. If you take ordinary water and, gradually cooling it, follow the change in density, you will notice that at the beginning a completely normal and natural process will occur - the water becomes denser and denser as it cools, and we will not see any deviations from the norm until until the water cools down to 4°C. Below this temperature, in spite of general ideas water suddenly becomes lighter, and when it freezes it becomes even lighter and forms ice that floats on the surface of the water. Freezing, water expands by 9% in relation to the previous volume. This expansion can be fatal for the water supply in the event of unexpected frosts. Water freezing in pipes will break them.

It is this feature of water, as is known, that protects lakes and ponds from continuous freezing in severe winters and thereby saves life in these reservoirs. Autumn air cools the surface layers of the lake, they become heavier and sink to the bottom. The lake is cooling. But this process only goes on until the water temperature reaches 4 °C. If now the surface layers become even colder, then they no longer sink to the bottom, since the density of these layers is less than the density of deep water, where a temperature of 4 ° C is maintained. Differences in density are not large - these differences appear only in the fourth decimal place - but these differences are quite enough so that water with a temperature close to 0 ° C cannot penetrate into the depths of the lake. Cooling process surface layers it will go faster now and soon the lead surface of the lake will be covered with the first brittle ice. Ice is a poor conductor of heat, it will reliably hide the life of the lake from terrible winter frosts. This circulation explains why ice forms earlier in shallower parts of the lake and thicker later.

The difference in temperatures of the upper and lower layers of water is used when dredgers operate in winter conditions. With the help of pumps from the deeper part of the reservoir, water is pumped into the surface layers, which prevents the formation of ice in the operating unit.

But sea ​​water(which, as you know, is a brine, each liter of which contains about 35 grams of salts) behaves completely differently when cooled: highest density it is observed at lower temperatures than fresh water, namely at -3.5 ° C. But sea water freezes at -1.9 ° C, i.e. it turns into ice before reaching its maximum density.

If, during the melting of ice, the volume of the liquid obtained is less than the volume of the ice taken, then it can be assumed that the transition of ice into liquid state will be lightened if the ice is subjected to pressure, i.e. bringing the crystals closer together. In fact, if high pressure is applied to ice, its melting point decreases. So, under a pressure of 2045 atm (per 1 cm 2), ice will melt at a temperature of -22 ° C. A further increase in pressure no longer reduces the melting point, since new forms of ice with new properties are formed. The ability of ice to melt at a lower temperature great pressure it also explains that in glaciers, whose thickness is enormous, melting at the base begins earlier than on the surface.

heat capacity of water

The amount of heat required to heat 1 g of water by 1° is enough to heat 9.25 g of iron, 10.3 g of copper by 1°. The abnormally high heat capacity of water turns the seas and oceans into a giant thermostat, smoothing out daily fluctuations in air temperature. Moreover, not only large masses of water, like the seas, are ways to smooth out these fluctuations, but also the usual water vapor of the atmosphere. Sharp diurnal fluctuations in temperature in the regions of the great deserts are associated with the absence of water vapor in the air. The dry air of the desert is almost devoid of water vapor, which could contain the rapid night cooling of the sand that has heated up during the day, so the air temperature can be no more than 5 ° C.

The heat capacity of water explains the phenomenon of different heating of water and land: since the heat capacity of the solid rocks that make up the land surface and the heat capacity of water differ sharply, different amounts of heat will be required to heat water and sand to the same temperature, so during the day the temperature of the sand is higher than water. Water cools more slowly than solid rock, so sand is colder at night than water. As you know, air is heated not directly by the rays of the sun, but by heat transfer from the heated surface of land and water. AT summer time a significant temperature difference is created between the surface of land and water, due to which air moves in the direction determined by the temperature difference between the waters of the seas and oceans and the land adjacent to them.

The heat capacity of water (1 cal), by the way, is 2 times greater than the heat capacity of ice (0.5 cal), and for all other substances, melting has almost no effect on this value.

Why does this value show such a large value in the case of water? Specific heat capacity is the amount of heat that must be imparted to one gram of a substance in order to increase its temperature by one degree Celsius. Consequently, water requires an abnormally large amount of heat for its heating. Since an increase in temperature means an increase in the average speed of movement of molecules, in molecular language, the high heat capacity of water means that its molecules are very inert. To enlarge average speed H 2 O molecules, for some reason they need to impart quite a lot of energy, although the molecules themselves are relatively small in terms of molecular scales. Everything is explained by the existence of hydrogen bonds. Since most of the molecules are bound into rather large complexes, a separate "average" H 2 O molecule can increase its kinetic energy in one of two ways. It can, firstly, having freed itself from all its hydrogen bonds, begin to move independently. And secondly, the acceleration of the entire complex of molecules will, of course, lead to an increase in the speed of each H 2 O molecule included in this complex. Obviously, both of these methods require significant energy costs, which leads to great importance specific heat water.

· Latent heat melting and evaporating water

If the temperature of the solid has risen to the melting point or if the liquid has reached the boiling point, then a transitional phase sets in, as if a pause, during which two phases (solid and liquid or liquid and gaseous) exist simultaneously. During this period of time, which lasts until the solid body is completely liquefied or liquefied to vapor, the absorbed heat does not cause any change in the temperature of the body. This heat is called latent heat, and its amount varies from substance to substance. The latent heat of fusion, as well as evaporation, is unusually high in water; this circumstance is of great importance for the temperature of the earth's surface. The word "hidden" we use already contains some allusion to one physical law that needs to be emphasized: the heat absorbed by water does not disappear anywhere. As you know, one of the basic laws of nature is the law of conservation and transformation of energy. In the very general view this law is formulated as follows: energy from one form passes into another (for example, thermal energy can turn into mechanical) without being destroyed; in closed system the total amount of energy remains constant. This law is also confirmed by the case we have cited. When we say that water has an exceptional heat capacity, we are simply stating that water as a substance can store more thermal energy with less movement of atoms and molecules (and this is exactly what is measured by temperature) than any other widespread substance. The energy stays in place, in the water; it will be released as heat when the ambient temperature drops; as a result, the decrease in temperature will not be so sharp. When water freezes, it gives off the same amount of heat that it absorbs when ice melts. We know that it is more difficult to endure hot, but damp weather with a temperature of about 30 °, than dry and clear weather with even more high temperature. The reason for this is twofold: firstly, our sweat, evaporating, cools us, taking away heat from the surface of the skin and from the surrounding air, but it cannot evaporate in the atmosphere of a damp day saturated with water vapor; secondly, when water vapor condenses and turns into water, exactly as much heat is released as it was spent on evaporation.

Water has the highest latent heat of vaporization and latent heat of fusion in the world of minerals. It takes five and a half times more heat to boil water out of a kettle than to boil it. If it were not for this property - even in the heat to slowly evaporate, many lakes and rivers would dry up to the bottom in summer. It takes a lot of heat to melt ice. The latent heat of fusion (the amount of heat required to melt 1 g of ice at 0°C) is 79.4 cal. This is why the spring melting of ice is slow and saves us from big floods (although not always).

Dielectric constant of water

Main electrical characteristic of any medium - the dielectric constant - in the case of water demonstrates features unusual for a liquid. Firstly, it is very large, for static electric fields it is 81, while for most other substances it does not exceed 10. If any substance is exposed to an alternating electric field, then the permittivity will cease to be constant value, but depends on the frequency of the applied field, decreasing strongly for high-frequency fields. But the permittivity of water decreases not only in time-varying fields, but also in space. variable fields, i.e. water is a nonlocally polarizable medium.

Great importance permittivity due to the peculiarities of the H 2 O molecule. The large value of the static permittivity of water ε = 81 is due to the fact that water is a highly polar liquid and therefore has a soft orientational degree of freedom (ie, rotation of molecular dipoles). Each water molecule has a significant dipole moment. In the absence of an electric field, the dipoles are randomly oriented, and the total electric field created by them is zero. If water is placed in an electric field, then the dipoles will begin to reorient so as to weaken the applied field. Such a picture is also observed in any other polar liquid, but due to the large value of the dipole moment of H 2 O molecules, water can very strongly (80 times) weaken the external field. This is how water reacts to an external electric field if the applied field is constant in time and changes little (or does not change at all) in the space filled with water. In variables electric fields the dielectric constant of water decreases with increasing frequency of the applied field, reaching a value of 4-5 for frequencies greater than 10 12 Hz. In 1929, P. Debye proposed to describe the reaction of water to an external electric field using the complex permittivity:

ε(ω) = ε ∞ + (ε ο - ε ∞)/(1 + i ω τ)

where ω is the frequency of the external electric field, i is the imaginary unit, τ is the characteristic relaxation time, ε ∞ ≈ 4÷5 is the permittivity of water at the highest frequency of the external field.

Although Debye used a rather artificial model of the structure of water in deriving his formula, this expression is in good agreement with experimental data. As we can see, as the frequency of the external field increases, the dielectric permittivity drops sharply. The molecular explanation for this phenomenon is quite simple. Any individual movement of the H 2 O molecule is strongly limited by hydrogen bonds. In alternating electric fields, molecular dipoles tend to follow the changing field. At low frequencies, they succeed. However, as the frequency increases, it becomes more and more difficult to navigate. Eventually, the dipoles stop responding to the external field altogether. The permittivity is now determined only by a fast atomic-molecular redistribution mechanism electric charge, which is inherent in all substances. Such mechanisms operate in water also in the case of constant fields, but their contribution to overall value dielectric constant is small, only 4-5 units.

Surface tension of water

You see its manifestation whenever you watch water slowly dripping from a faucet. A film of water emerges from the faucet and begins to stretch, like a thin rubber shell, under the weight of the liquid contained in it. This film, attached to the faucet opening, gradually lengthens until its weight suddenly becomes too great. The film, however, does not break, as a cutter would break if overloaded. Instead, it "slides" off the coccyx of the faucet and, as if embracing a small amount of water, forms a freely falling droplet. Undoubtedly, you have observed more than once that falling droplets take on an almost spherical shape. If there were no external forces, they would be strictly spherical. What you are observing is one of the manifestations of the unusual ability of water to "contract", "self-compact", or, in other words, its ability to cohesion (cohesion). A drop of water dripping from a faucet shrinks into a tiny ball, and a ball of all possible geometric bodies has the smallest surface area for a given volume.

Due to adhesion, tension is formed on the surface of the water, and in order to break the surface of the water, it takes physical strength, and, oddly enough, quite significant. An undisturbed water surface can hold objects that are much "heavier" than water, such as a steel needle or a razor blade, or some insects that glide through the water as if it were not a liquid, but a solid body.

Of all liquids except mercury, water has the highest surface tension.

Inside the liquid, the attraction of molecules to each other is balanced. But not on the surface. Water molecules that lie deeper pull down the topmost molecules. Therefore, a drop of water, as it were, tends to shrink as much as possible. It is pulled together by surface tension forces.

Physicists calculated exactly which weight should be hung from a column of water three centimeters thick in order to break it. The weight will need a huge one - more than a hundred tons! But this is when the water is exceptionally clean. There is no such water in nature. There is always something in it. Let at least a little, but foreign substances break the links in the strong chain of water molecules, and the cohesive forces between them decrease.

If drops of mercury are applied to a glass plate, and drops of water to a paraffin one, then very small droplets will have the shape of a ball, while larger ones will be slightly flattened by gravity.

This phenomenon is explained by the fact that between mercury and glass, as well as between paraffin and water, attractive forces (adhesion) arise that are smaller than between the molecules themselves (cohesion). When water comes into contact with clean glass, and mercury with metal plate we observe an almost uniform distribution of both substances on the plates, since the forces of attraction between glass and water molecules, metal and mercury molecules are greater than the attraction between individual molecules of water and mercury. This phenomenon, when the liquid is evenly distributed on the surface solid body, is called wetting. This means that water wets clean glass, but does not wet paraffin. Wettability in a particular case can indicate the degree of contamination of the surface. For example, on a cleanly washed plate (porcelain, earthenware), water spreads evenly, in a cleanly washed flask the walls are evenly covered with water, but if the water on the surface takes the form of drops, this indicates that the surface of the dish is covered with a thin layer of a substance that does not wet water , most often fat.

Examples of the structure of water:

1. A crystal of distilled water, unaffected.

2. Spring water.

3. Antarctic ice.

4. This is what a water crystal looks like after listening to Beethoven's Pastoral.

5. A crystal formed after listening to heavy metal rock.

6. The crystal after exposure to the words "You are a fool" is very similar to the crystal after exposure to "hard rock" music.

7. The word "Angel".

8. The word "Devil".

9. Water received a request to "Do it."

10. Water received the order "Do it."

11. The words “You bored me. I'll kill you".

12. Water received electromagnetic radiation love and gratitude

17. The words "Love and Gratitude" spoken in English.

18. The words "Love and Gratitude" spoken in Japanese.

19. The words "Love and gratitude" spoken on German

Amphiphilic substances:

Everyone knows that fish feel comfortable only in the water, and most cats treat water procedures with obvious discontent, but animals such as a frog or a newt are quite capable of both swimming in a river or a puddle and moving freely on the ground! These animals are called amphibians or amphibians. Their amphibians, capable of dissolving both in hydrophilic and hydrophobic, the molecules of amphiphilic compounds themselves are similar to a tadpole: they consist of a long hydrocarbon tail (usually built from more than ten CH 2 groups), which ensures solubility in non-polar media, and a polar head, responsible for hydrophilic properties. Thus, amphiphilic compounds simultaneously “love” both water (that is, they are hydrophilic) and non-polar solvents (show hydrophobic properties).

Depending on the type of hydrophilic group, amphiphilic compounds bearing a charged cationic or anionic functional group and amphiphilic compounds with an uncharged functional group are isolated. The vast majority of known organic compounds carry more than one charged functional group. An example of such substances are macromolecular compounds - proteins, lipoproteins, block copolymers, etc. The presence of a tertiary structure in protein molecules, which is formed as a result of intramolecular interactions of functional groups (polar or nonpolar) with each other, in itself demonstrates the amphiphilic nature of these compounds.

Another example of amphiphilic compounds is most drugs, the molecules of which combine a set of certain functional groups necessary for effective binding to the target receptor.

The role of amphiphilic compounds in obtaining nanomaterials and nanotechnological products is difficult to overestimate. Amphiphilic compounds are often surfactants. Their molecules “self-assemble” (self-assemble) at different interfaces, forming thin films of self-assembling monolayers with a thickness of only one molecule, forming “micellar” systems.

Amphiphilic compounds play special role in living nature. No animal or plant can exist without them! It is from amphiphilic molecules that the cell membrane consists, which separates a living organism from a hostile one. external environment. It is these molecules that make up the internal organelles of the cell, participate in the process of its division, are involved in the metabolism with environment. Amphiphilic molecules serve as food for us and are formed in our bodies, participate in the internal regulation and cycle of bile acids. Our body contains more than 10% amphiphilic molecules. That is why synthetic surfactants can be dangerous for living organisms and, for example, can dissolve the cell membrane and lead to its death.