All chemical properties. Chemical, physical properties of substances

Group IIA contains only metals – Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium) and Ra (radium). The chemical properties of the first representative of this group, beryllium, differ most strongly from the chemical properties of the other elements of this group. Its chemical properties are in many ways even more similar to aluminum than to other Group IIA metals (so-called “diagonal similarity”). Magnesium, in its chemical properties, also differs markedly from Ca, Sr, Ba and Ra, but still has much more similar chemical properties with them than with beryllium. Due to the significant similarity in the chemical properties of calcium, strontium, barium and radium, they are combined into one family called alkaline earth metals.

All elements of group IIA belong to s-elements, i.e. contain all their valence electrons on s-sublevel Thus, the electronic configuration of the outer electronic layer of all chemical elements of this group has the form ns 2 , Where n– number of the period in which the element is located.

Due to the peculiarities of the electronic structure of group IIA metals, these elements, in addition to zero, can have only one single oxidation state equal to +2. Simple substances formed by elements of group IIA, when participating in any chemical reactions, are only capable of oxidation, i.e. donate electrons:

Me 0 – 2e — → Me +2

Calcium, strontium, barium and radium have extremely high chemical reactivity. The simple substances formed by them are very strong reducing agents. Magnesium is also a strong reducing agent. The reduction activity of metals obeys the general laws of the periodic law of D.I. Mendeleev and increases down the subgroup.

Interaction with simple substances

with oxygen

Without heating, beryllium and magnesium do not react with either atmospheric oxygen or pure oxygen due to the fact that they are covered with thin protective films consisting of BeO and MgO oxides, respectively. Their storage does not require any special methods of protection from air and moisture, unlike alkaline earth metals, which are stored under a layer of liquid inert to them, most often kerosene.

Be, Mg, Ca, Sr, when burned in oxygen, form oxides of the composition MeO, and Ba - a mixture of barium oxide (BaO) and barium peroxide (BaO 2):

2Mg + O2 = 2MgO

2Ca + O2 = 2CaO

2Ba + O 2 = 2BaO

Ba + O 2 = BaO 2

It should be noted that when alkaline earth metals and magnesium burn in air, a side reaction of these metals with air nitrogen also occurs, as a result of which, in addition to compounds of metals with oxygen, nitrides with the general formula Me 3 N 2 are also formed.

with halogens

Beryllium reacts with halogens only at high temperatures, and the rest of the Group IIA metals - already at room temperature:

Mg + I 2 = MgI 2 – Magnesium iodide

Ca + Br 2 = CaBr 2 – calcium bromide

Ba + Cl 2 = BaCl 2 – barium chloride

with non-metals of groups IV–VI

All metals of group IIA react when heated with all nonmetals of groups IV–VI, but depending on the position of the metal in the group, as well as the activity of the nonmetals, varying degrees of heating are required. Since beryllium is the most chemically inert among all group IIA metals, when carrying out its reactions with non-metals, significant use is required. O higher temperature.

It should be noted that the reaction of metals with carbon can form carbides of different natures. There are carbides that belong to methanides and are conventionally considered derivatives of methane, in which all hydrogen atoms are replaced by metal. They, like methane, contain carbon in the -4 oxidation state, and when they are hydrolyzed or interact with non-oxidizing acids, one of the products is methane. There is also another type of carbides - acetylenides, which contain the C 2 2- ion, which is actually a fragment of the acetylene molecule. Carbides such as acetylenides, upon hydrolysis or interaction with non-oxidizing acids, form acetylene as one of the reaction products. The type of carbide - methanide or acetylenide - obtained when a particular metal reacts with carbon depends on the size of the metal cation. Metal ions with a small radius usually form metanides, and larger ions form acetylenides. In the case of metals of the second group, methanide is obtained by the interaction of beryllium with carbon:

The remaining metals of group II A form acetylenides with carbon:

With silicon, group IIA metals form silicides - compounds of the type Me 2 Si, with nitrogen - nitrides (Me 3 N 2), with phosphorus - phosphides (Me 3 P 2):

with hydrogen

All alkaline earth metals react with hydrogen when heated. In order for magnesium to react with hydrogen, heating alone, as in the case of alkaline earth metals, is not enough; in addition to high temperature, increased hydrogen pressure is also required. Beryllium does not react with hydrogen under any conditions.

Interaction with complex substances

with water

All alkaline earth metals react actively with water to form alkalis (soluble metal hydroxides) and hydrogen. Magnesium reacts with water only when boiled due to the fact that when heated, the protective oxide film MgO dissolves in water. In the case of beryllium, the protective oxide film is very resistant: water does not react with it either when boiling or even at red-hot temperatures:

with non-oxidizing acids

All metals of the main subgroup of group II react with non-oxidizing acids, since they are in the activity series to the left of hydrogen. In this case, a salt of the corresponding acid and hydrogen are formed. Examples of reactions:

Be + H 2 SO 4 (diluted) = BeSO 4 + H 2

Mg + 2HBr = MgBr 2 + H 2

Ca + 2CH 3 COOH = (CH 3 COO) 2 Ca + H 2

with oxidizing acids

− diluted nitric acid

All metals of group IIA react with dilute nitric acid. In this case, the reduction products, instead of hydrogen (as in the case of non-oxidizing acids), are nitrogen oxides, mainly nitrogen oxide (I) (N 2 O), and in the case of highly dilute nitric acid, ammonium nitrate (NH 4 NO 3):

4Ca + 10HNO3 ( razb .) = 4Ca(NO 3) 2 + N 2 O + 5H 2 O

4Mg + 10HNO3 (very blurry)= 4Mg(NO 3) 2 + NH 4 NO 3 + 3H 2 O

− concentrated nitric acid

Concentrated nitric acid at ordinary (or low) temperature passivates beryllium, i.e. does not react with it. When boiling, the reaction is possible and proceeds predominantly in accordance with the equation:

Magnesium and alkaline earth metals react with concentrated nitric acid to form a wide range of different nitrogen reduction products.

− concentrated sulfuric acid

Beryllium is passivated with concentrated sulfuric acid, i.e. does not react with it under normal conditions, but the reaction occurs at boiling and leads to the formation of beryllium sulfate, sulfur dioxide and water:

Be + 2H 2 SO 4 → BeSO 4 + SO 2 + 2H 2 O

Barium is also passivated by concentrated sulfuric acid due to the formation of insoluble barium sulfate, but reacts with it when heated; barium sulfate dissolves when heated in concentrated sulfuric acid due to its conversion to barium hydrogen sulfate.

The remaining metals of main group IIA react with concentrated sulfuric acid under any conditions, including in the cold. Reduction of sulfur can occur to SO 2, H 2 S and S depending on the activity of the metal, reaction temperature and acid concentration:

Mg + H2SO4 ( conc. .) = MgSO 4 + SO 2 + H 2 O

3Mg + 4H 2 SO 4 ( conc. .) = 3MgSO 4 + S↓ + 4H 2 O

4Ca + 5H 2 SO 4 ( conc. .) = 4CaSO 4 +H 2 S + 4H 2 O

with alkalis

Magnesium and alkaline earth metals do not interact with alkalis, and beryllium easily reacts both with alkali solutions and with anhydrous alkalis during fusion. Moreover, when a reaction is carried out in an aqueous solution, water also participates in the reaction, and the products are tetrahydroxoberyllates of alkali or alkaline earth metals and hydrogen gas:

Be + 2KOH + 2H 2 O = H 2 + K 2 - potassium tetrahydroxoberyllate

When carrying out a reaction with a solid alkali during fusion, beryllates of alkali or alkaline earth metals and hydrogen are formed

Be + 2KOH = H 2 + K 2 BeO 2 - potassium beryllate

with oxides

Alkaline earth metals, as well as magnesium, can reduce less active metals and some nonmetals from their oxides when heated, for example:

The method of reducing metals from their oxides with magnesium is called magnesium.

If in D.I. Mendeleev’s periodic table of elements we draw a diagonal from beryllium to astatine, then on the lower left along the diagonal there will be metal elements (these also include elements of side subgroups, highlighted in blue), and on the upper right - non-metal elements (highlighted yellow). Elements located near the diagonal - semimetals or metalloids (B, Si, Ge, Sb, etc.) have a dual character (highlighted in pink).

As can be seen from the figure, the vast majority of elements are metals.

By their chemical nature, metals are chemical elements whose atoms give up electrons from external or pre-external energy levels, forming positively charged ions.

Almost all metals have relatively large radii and a small number of electrons (from 1 to 3) at the outer energy level. Metals are characterized by low electronegativity values ​​and reducing properties.

The most typical metals are located at the beginning of the periods (starting from the second), then from left to right the metallic properties weaken. In the group from top to bottom, the metallic properties increase as the radius of the atoms increases (due to an increase in the number of energy levels). This leads to a decrease in electronegativity (the ability to attract electrons) of elements and an increase in reducing properties (the ability to donate electrons to other atoms in chemical reactions).

Typical metals are s-elements (elements of the IA group from Li to Fr. elements of the PA group from Mg to Ra). The general electronic formula of their atoms is ns 1-2. They are characterized by oxidation states + I and + II, respectively.

The small number of electrons (1-2) in the outer energy level of typical metal atoms means that these electrons are easily lost and exhibit strong reducing properties, as reflected by low electronegativity values. This implies the limited chemical properties and methods of obtaining typical metals.

A characteristic feature of typical metals is the tendency of their atoms to form cations and ionic chemical bonds with non-metal atoms. Compounds of typical metals with nonmetals are ionic crystals of “metalanion of nonmetal,” for example K + Br -, Ca 2+ O 2-. Cations of typical metals are also included in compounds with complex anions - hydroxides and salts, for example Mg 2+ (OH -) 2, (Li +)2CO 3 2-.

The A-group metals that form the amphoteric diagonal in the Periodic Table Be-Al-Ge-Sb-Po, as well as the metals adjacent to them (Ga, In, Tl, Sn, Pb, Bi) do not exhibit typical metallic properties. General electronic formula of their atoms ns 2 n.p. 0-4 involves a greater variety of oxidation states, a greater ability to retain their own electrons, a gradual decrease in their reducing ability and the appearance of oxidizing ability, especially in high oxidation states (typical examples are compounds Tl III, Pb IV, Bi v). Similar chemical behavior is characteristic of most (d-elements, i.e. elements of the B-groups of the Periodic Table (typical examples are the amphoteric elements Cr and Zn).

This manifestation of duality (amphoteric) properties, both metallic (basic) and non-metallic, is due to the nature of the chemical bond. In the solid state, compounds of atypical metals with nonmetals contain predominantly covalent bonds (but less strong than bonds between nonmetals). In solution, these bonds are easily broken, and the compounds dissociate into ions (in whole or in part). For example, the metal gallium consists of Ga 2 molecules; in the solid state, the chlorides of aluminum and mercury (II) AlCl 3 and HgCl 2 contain strongly covalent bonds, but in solution AlCl 3 dissociates almost completely, and HgCl 2 - to a very small extent (and then into HgCl + and Cl - ions).

General physical properties of metals

Due to the presence of free electrons ("electron gas") in the crystal lattice, all metals exhibit the following characteristic general properties:

1) Plastic- the ability to easily change shape, stretch into wire, and roll into thin sheets.

2) Metallic shine and opacity. This is due to the interaction of free electrons with light incident on the metal.

3) Electrical conductivity. It is explained by the directional movement of free electrons from the negative pole to the positive one under the influence of a small potential difference. When heated, electrical conductivity decreases, because As the temperature increases, vibrations of atoms and ions in the nodes of the crystal lattice intensify, which complicates the directional movement of the “electron gas”.

4) Thermal conductivity. It is caused by the high mobility of free electrons, due to which the temperature quickly equalizes over the mass of the metal. The highest thermal conductivity is found in bismuth and mercury.

5) Hardness. The hardest is chrome (cuts glass); the softest alkali metals - potassium, sodium, rubidium and cesium - are cut with a knife.

6) Density. The smaller the atomic mass of the metal and the larger the radius of the atom, the smaller it is. The lightest is lithium (ρ=0.53 g/cm3); the heaviest is osmium (ρ=22.6 g/cm3). Metals with a density of less than 5 g/cm3 are considered “light metals”.

7) Melting and boiling points. The most fusible metal is mercury (mp = -39°C), the most refractory metal is tungsten (mp = 3390°C). Metals with melting temperature above 1000°C are considered refractory, below – low-melting.

General chemical properties of metals

Strong reducing agents: Me 0 – nē → Me n +

A number of voltages characterize the comparative activity of metals in redox reactions in aqueous solutions.

I. Reactions of metals with non-metals

1) With oxygen:
2Mg + O 2 → 2MgO

2) With sulfur:
Hg + S → HgS

3) With halogens:
Ni + Cl 2 – t° → NiCl 2

4) With nitrogen:
3Ca + N 2 – t° → Ca 3 N 2

5) With phosphorus:
3Ca + 2P – t° → Ca 3 P 2

6) With hydrogen (only alkali and alkaline earth metals react):
2Li + H 2 → 2LiH

Ca + H 2 → CaH 2

II. Reactions of metals with acids

1) Metals in the electrochemical voltage series up to H reduce non-oxidizing acids to hydrogen:

Mg + 2HCl → MgCl 2 + H 2

2Al+ 6HCl → 2AlCl 3 + 3H 2

6Na + 2H 3 PO 4 → 2Na 3 PO 4 + 3H 2

2) With oxidizing acids:

When nitric acid of any concentration and concentrated sulfuric acid interact with metals Hydrogen is never released!

Zn + 2H 2 SO 4(K) → ZnSO 4 + SO 2 + 2H 2 O

4Zn + 5H 2 SO 4(K) → 4ZnSO 4 + H 2 S + 4H 2 O

3Zn + 4H 2 SO 4(K) → 3ZnSO 4 + S + 4H 2 O

2H 2 SO 4 (k) + Cu → Cu SO 4 + SO 2 + 2H 2 O

10HNO 3 + 4Mg → 4Mg(NO 3) 2 + NH 4 NO 3 + 3H 2 O

4HNO 3 (k) + Cu → Cu (NO 3) 2 + 2NO 2 + 2H 2 O

III. Interaction of metals with water

1) Active (alkali and alkaline earth metals) form a soluble base (alkali) and hydrogen:

2Na + 2H 2 O → 2NaOH + H 2

Ca+ 2H 2 O → Ca(OH) 2 + H 2

2) Metals of medium activity are oxidized by water when heated to an oxide:

Zn + H 2 O – t° → ZnO + H 2

3) Inactive (Au, Ag, Pt) - do not react.

IV. Displacement of less active metals by more active metals from solutions of their salts:

Cu + HgCl 2 → Hg+ CuCl 2

Fe+ CuSO 4 → Cu+ FeSO 4

In industry, they often use not pure metals, but mixtures of them - alloys, in which the beneficial properties of one metal are complemented by the beneficial properties of another. Thus, copper has low hardness and is unsuitable for the manufacture of machine parts, while alloys of copper and zinc ( brass) are already quite hard and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but is too soft. Based on it, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing the beneficial properties of aluminum, acquires high hardness and becomes suitable for aircraft construction. Alloys of iron with carbon (and additives of other metals) are widely known cast iron And steel.

Free metals are restorers. However, some metals have low reactivity due to the fact that they are coated surface oxide film, to varying degrees, resistant to chemical reagents such as water, solutions of acids and alkalis.

For example, lead is always covered with an oxide film; its transition into solution requires not only exposure to a reagent (for example, dilute nitric acid), but also heating. The oxide film on aluminum prevents its reaction with water, but is destroyed by acids and alkalis. Loose oxide film (rust), formed on the surface of iron in moist air, does not interfere with further oxidation of iron.

Under the influence concentrated acids form on metals sustainable oxide film. This phenomenon is called passivation. So, in concentrated sulfuric acid metals such as Be, Bi, Co, Fe, Mg and Nb are passivated (and then do not react with acid), and in concentrated nitric acid - metals A1, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb , Th and U.

When interacting with oxidizing agents in acidic solutions, most metals transform into cations, the charge of which is determined by the stable oxidation state of a given element in compounds (Na +, Ca 2+, A1 3+, Fe 2+ and Fe 3+)

The reducing activity of metals in an acidic solution is transmitted by a series of stresses. Most metals are transferred into solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only with sulfuric (concentrated) and nitric acids, and Pt and Au - with “regia vodka”.

Metal corrosion

An undesirable chemical property of metals is their active destruction (oxidation) upon contact with water and under the influence of oxygen dissolved in it (oxygen corrosion). For example, the corrosion of iron products in water is widely known, as a result of which rust forms and the products crumble into powder.

Corrosion of metals also occurs in water due to the presence of dissolved gases CO 2 and SO 2; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 ( hydrogen corrosion).

The area of ​​contact between two dissimilar metals can be especially corrosive ( contact corrosion). A galvanic couple occurs between one metal, for example Fe, and another metal, for example Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the voltage series (Re), to the less active metal (Sn, Cu), and the more active metal is destroyed (corroded).

It is because of this that the tinned surface of cans (iron coated with tin) rusts when stored in a humid atmosphere and handled carelessly (iron quickly collapses after even a small scratch appears, allowing the iron to come into contact with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, since even if there are scratches, it is not the iron that corrodes, but the zinc (a more active metal than iron).

Corrosion resistance for a given metal increases when it is coated with a more active metal or when they are fused; Thus, coating iron with chromium or making an alloy of iron and chromium eliminates corrosion of iron. Chromed iron and steel containing chromium ( stainless steel), have high corrosion resistance.

electrometallurgy, i.e., obtaining metals by electrolysis of melts (for the most active metals) or salt solutions;

pyrometallurgy, i.e., the recovery of metals from ores at high temperatures (for example, the production of iron in the blast furnace process);

hydrometallurgy, i.e., the separation of metals from solutions of their salts by more active metals (for example, the production of copper from a solution of CuSO 4 by the action of zinc, iron or aluminum).

Native metals are sometimes found in nature (typical examples are Ag, Au, Pt, Hg), but more often metals are found in the form of compounds ( metal ores). Metals vary in abundance in the earth's crust: from the most common - Al, Na, Ca, Fe, Mg, K, Ti) to the rarest - Bi, In, Ag, Au, Pt, Re.

The chemical properties of a substance depend not only on what chemical elements it consists of, but also on the structure of the molecules of the substance (structural isomerism) and on the spatial configuration of the molecules (conformation, stereoisomerism). As a rule, substances that have the same composition and structure have the same chemical properties, with the exception of reactions with substances of a different spatial configuration. This distinction is especially important in biochemistry, for example, the ability of a protein to react with other biologically active substances may depend on the way it folds.

Examples of chemical properties

see also

Notes

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For the last 200 years of humanity studied the properties of substances better than in the entire history of the development of chemistry. Naturally, the number of substances is also growing rapidly; this is due, first of all, to the development of various methods for obtaining substances.

In everyday life we ​​come across many substances. Among them are water, iron, aluminum, plastic, soda, salt and many others. Substances that exist in nature, such as oxygen and nitrogen contained in the air, substances dissolved in water and of natural origin, are called natural substances. Aluminum, zinc, acetone, lime, soap, aspirin, polyethylene and many other substances do not exist in nature.

They are obtained in the laboratory and produced by industry. Artificial substances are not found in nature; they are created from natural substances. Some substances that exist in nature can also be obtained in a chemical laboratory.

Thus, when potassium permanganate is heated, oxygen is released, and when chalk is heated, oxygen is released. carbon dioxide. Scientists have learned to turn graphite into diamond; they are growing crystals of ruby, sapphire and malachite. So, along with substances of natural origin, there are a huge number of artificially created substances that are not found in nature.

Substances not found in nature are produced in various enterprises: factories, factories, combines, etc.

In the context of depletion of the natural resources of our planet, chemists now face an important task: to develop and implement methods by which it is possible to artificially, in a laboratory or industrial production, obtain substances that are analogues of natural substances. For example, reserves of fossil fuels in nature are running out.

There may come a time when oil and natural gas run out. Already, new types of fuel are being developed that would be just as efficient, but would not pollute the environment. Today, humanity has learned to artificially obtain various precious stones, for example, diamonds, emeralds, and beryls.

State of matter

Substances can exist in several states of aggregation, three of which are known to you: solid, liquid, gaseous. For example, water in nature exists in all three states of aggregation: solid (in the form of ice and snow), liquid (liquid water) and gaseous (water vapor). There are known substances that cannot exist under normal conditions in all three states of aggregation. For example, such a substance is carbon dioxide. At room temperature it is an odorless and colorless gas. At a temperature of –79°C this substance “freezes” and turns into a solid state of aggregation. The everyday (trivial) name for such a substance is “dry ice”. This name is given to this substance due to the fact that “dry ice” turns into carbon dioxide without melting, that is, without transitioning to a liquid state of aggregation, which is present, for example, in water.

Thus, an important conclusion can be drawn. A substance, when transitioning from one state of aggregation to another, does not transform into other substances. The process of a certain change, transformation, is called a phenomenon.

Physical phenomena. Physical properties of substances.

Phenomena in which substances change their state of aggregation, but do not transform into other substances, are called physical. Each individual substance has certain properties. The properties of substances may be different or similar to each other. Each substance is described using a set of physical and chemical properties. Let's take water as an example. Water freezes and turns into ice at a temperature of 0°C, and boils and turns into steam at a temperature of +100°C. These phenomena are considered physical, since water has not turned into other substances, only a change in the state of aggregation occurs. These freezing and boiling points are physical properties specific to water.

Properties of substances that are determined by measurements or visually in the absence of transformation of some substances into others are called physical

Evaporation of alcohol, like evaporation of water– physical phenomena, substances in this case change their state of aggregation. After the experiment, you can be sure that alcohol evaporates faster than water - these are the physical properties of these substances.

The main physical properties of substances include the following: state of aggregation, color, odor, solubility in water, density, boiling point, melting point, thermal conductivity, electrical conductivity. Physical properties such as color, smell, taste, crystal shape can be determined visually using the senses, and density, electrical conductivity, melting and boiling points are determined by measurement. Information about the physical properties of many substances is collected in specialized literature, for example, in reference books. The physical properties of a substance depend on its state of aggregation. For example, the densities of ice, water and water vapor are different.

Gaseous oxygen is colorless, but liquid oxygen is blue. Knowledge of physical properties helps to “recognize” many substances. For example, copper- The only metal that is red in color. Only table salt has a salty taste. Iodine- An almost black solid that turns into a purple vapor when heated. In most cases, to identify a substance, you need to consider several of its properties. As an example, let us characterize the physical properties of water:

• color – colorless (in small volumes)
• smell - no smell
• state of aggregation - liquid under normal conditions
• density – 1 g/ml,
• boiling point – +100°С
• melting point – 0°C
• thermal conductivity – low
• electrical conductivity - pure water does not conduct electricity

Crystalline and amorphous substances

When describing the physical properties of solids, it is customary to describe the structure of the substance. If you examine a sample of table salt under a magnifying glass, you will notice that the salt consists of many tiny crystals. In salt deposits you can also find very large crystals. Crystals are solids in the shape of regular polyhedra. Crystals can have different shapes and sizes. Crystals of certain substances, such as table salt salt – fragile and easy to break. There are crystals that are quite hard. For example, diamond is considered one of the hardest minerals. If you examine table salt crystals under a microscope, you will notice that they all have a similar structure. If we consider, for example, glass particles, they will all have a different structure - such substances are called amorphous. Amorphous substances include glass, starch, amber, and beeswax. Amorphous substances are substances that do not have a crystalline structure

Chemical phenomena. Chemical reaction.

If during physical phenomena substances, as a rule, only change their state of aggregation, then during chemical phenomena the transformation of some substances into other substances occurs. Here are some simple examples: burning of a match is accompanied by charring of wood and the release of gaseous substances, that is, an irreversible transformation of wood into other substances occurs. Another example: Over time, bronze sculptures become covered with a green coating. The fact is that bronze contains copper. This metal slowly interacts with oxygen, carbon dioxide and air moisture, as a result of which new green substances are formed on the surface of the sculpture Chemical phenomena - phenomena of transformation of one substance into another The process of interaction of substances with the formation of new substances is called a chemical reaction. Chemical reactions occur all around us. Chemical reactions also occur within ourselves. In our body, transformations of many substances continuously occur; substances react with each other, forming reaction products. Thus, in a chemical reaction there are always reacting substances and substances formed as a result of the reaction.

• Chemical reaction– the process of interaction of substances, as a result of which new substances with new properties are formed
• Reagents- substances that enter into a chemical reaction
• Products– substances formed as a result of a chemical reaction

A chemical reaction is represented in general form by a reaction diagram REAGENTS -> PRODUCTS

• reagents– starting materials taken to carry out the reaction;
• products– new substances formed as a result of a reaction.

Any chemical phenomena (reactions) are accompanied by certain signs, with the help of which chemical phenomena can be distinguished from physical ones. Such signs include changes in the color of substances, the release of gas, the formation of sediment, the release of heat, and the emission of light.

Many chemical reactions are accompanied by the release of energy in the form of heat and light. As a rule, such phenomena are accompanied by combustion reactions. In combustion reactions in air, substances react with oxygen contained in the air. For example, the metal magnesium flares up and burns in air with a bright, blinding flame. This is why magnesium flash was used to create photographs in the first half of the 20th century. In some cases, it is possible to release energy in the form of light, but without releasing heat. One type of Pacific plankton is capable of emitting a bright blue light, clearly visible in the dark. The release of energy in the form of light is the result of a chemical reaction that occurs in the organisms of this type of plankton.

Summary of the article:

• There are two large groups of substances: substances of natural and artificial origin.
• Under normal conditions, substances can exist in three states of aggregation
• Properties of substances that are determined by measurements or visually in the absence of transformation of some substances into others are called physical
• Crystals are solids in the shape of regular polyhedra.
• Amorphous substances are substances that do not have a crystalline structure
• Chemical phenomena - phenomena of transformation of one substance into another
• Reagents are substances that enter into a chemical reaction.
• Products are substances formed as a result of a chemical reaction
• Chemical reactions can be accompanied by the release of gas, sediment, heat, light; change in color of substances
• Combustion is a complex physicochemical process of converting starting substances into combustion products during a chemical reaction, accompanied by intense release of heat and light (flame)

God gave man iron, but the devil gave him rust.

Proverb

Changes in properties over decades. Since d-elements are characterized by positive st.ok., then in the form of simple substances they exhibit reducing properties, which in aqueous solutions are characterized by the value of redox potential E. 0 In decades from left to right, its value, correlating with the value of I 1, growing, but when moving to manganese and the zinc subgroup, despite a sharp increase in I 1 , it decreases due to a decrease in the value of I 2 and a decrease in the energy of the crystal lattice when moving to these metals (from those located to the left of them in the periodic table).

In a compact state at rev. even M of the first decade, having negative values ​​E (0 from Sc to Mn E 0< −0,90 B), с водой не реагируют вследствие образованияpassivating oxide films on their surface. However, at red heat temperatures, less active metals (iron, nickel, analogues of vanadium and titanium) displace hydrogen from water. The reactivity of M also increases sharply when they are converted into fine state, for example, manganese and chromium powders interact with water at r.b. (with the formation of MnO 2 and Cr 2 O 3).

All metals of the first decade for which E 0 displace hydrogen from dilute acid solutions< 0, кроме ванадия. Наиболее активные М: цинк и марганец – растворяются даже в уксусной кислоте, а медь (в ряду напряжений стоит правее водорода) лишь в т.н. кислотах-окислителях. При указанных взаимодействиях только Sc и Тi образуют соединения в ст.ок. (+3), остальные – в (+2), хотя хром(II) и (гораздо медленнее) железо(II) на воздухе затем окисляются до (+3).

The anomalous passivity of vanadium (E 0 = −1.20 V) in dilute acids is explained special density its oxide film. It dissolves only in HF or concentrated HNO3, with which this metal reacts:

V + HNO 3 = HVO 3 + NO.

Other active M depending on solubility their oxide film in concentrated nitric acid either reacts with it, reducing nitrogen to (-3) (this is zinc, manganese and the scandium subgroup), or is passivated by it due to the thickening of the oxide film, such as Cr 124.

Passivation can also be carried out artificially. Thus, treating chromium (which is in the voltage range between zinc and iron) with concentrated nitric acid increases its potential from –0.56 V to +1.2 V, i.e. makes Cr almost as noble as Pt. (Chromium in stainless steel and other 125 alloys is especially easily deactivated.) Concentrated H 2 SO 4 and HNO 3 also passivate iron.

Cobalt and nickel are similar to Fe in chemical activity due to the proximity of atomic radii (therefore they are combined into familygland). However, if iron reacts with dilute HCl and H 2 SO 4 at ambient conditions, then Co and Ni react with heating. In addition, they are deactivated by nitric acid to a lesser extent than iron, due to the greater solubility of their oxides in this acid.

Note that for the elements of the second and third decades the nature of the change in the value of E 0 remains approximately the same as in the first.

Changes to properties in subgroups. The value of I 1 in d-subgroups is mainly grows andstrength increases bonds in the M lattice (compare m.p.). As a consequence (in contrast to the main subgroups and the Sc subgroup), the value of E 0 becomes more positive, and the reactivity of metals decreases.

So, in subgroup IB, if copper dissolves in concentrated sulfuric acid at r.v., then silver only at t > 160 0 C. However, silver, like copper, at roomtemperature interacts with nitric acid, and gold interacts only with aqua regia (as well as with selenic acid (see above) and with chlorine water in the presence of HCl).

In subgroup IIB, Zn is soluble even in acetic acid, Cd is soluble in HCl, and Hg (E 0 > 0) only in HNO (3 with a lack of acid, oxidation proceeds to Hg, 2 2 + and with an excess - to Hg). 2 +

Similarly, in subgroup VIIB - Mn reacts with CH COOH 3, and Tc and Re (their values

E 0: 0.47 V and 0.37 V, respectively), at r.b. dissolve only in oxidizing acids, for example, nitric acid (products NO and HEO 4).

In subgroup VIIIB, the metals of the iron family all interact with dilute acids. And their analogues, i.e. platinum metals (E 0 > 0) are oxidized only in tough conditions, and the proximity of their radii determines a large similarity in chemical behavior, but there is also differences.

Thus, the most active of them, palladium, is an acid, like silver; and rhodium and iridium, unlike the others, do not dissolve even in “regia vodka” 126. They react with a solution of sodium chloride saturated with chlorine at a red-hot temperature due to the formation sustainable complexes Na 3 [ECl 6 ]. However, in the form of black, these metals easily react with hot sulfuric acid and even with hydrochloric acid in the presence of oxygen. Note that under these conditions, osmium, due to its high affinity for oxygen (?), dissolves in a compact form.

In IV, V and VI side subgroups in M ​​of the second and third decades E 0< 0 , но за счет влиянияdense oxide film on their surface, they react with acids only under harsh conditions. Thus, Zr and Hf are soluble only in complexing acids: in hot sulfuric acid (product – H 2 [E(SO 4) ] 3) and in hydrofluoric acid (H 4 [EF 8 ]); molybdenum interacts only with oxidizing acids when heated, and tungsten, niobium and tantalum only with a mixture of HF and HNO (3 products NO and H 2 WF 8 or H 2 EF 7, respectively).

So, regardless of whether there is an imposition of a kinetic factor (passivating film) or not, the activity of d-metals towards acids in the subgroups decreases. Exception, as already noted, is scandium subgroup, in which there is no influence of f-compression and the nature of the change in the values ​​of the atomic radius, I 1 and E 0 is the same as in the main subgroups. As a consequence, lanthanum (unlike scandium and yttrium, which are soluble at r.b. only in acids) even interacts with water:

La + H 2 O → La(OH) 3 + H 2 .

Ratio of d-metals to alkalis. Silver 127 is most resistant to alkali, and zinc is the least resistant: even solution alkali, reducing the hydrogen of water and forming complex 128 -. The remaining d-metals, if they tend to exist in anionic form, react with alkalis (or soda) during fusion, For example:

Ti⎫ ⎧Na 2 TiO 3 ⎬ + NaOH→ H 2 + ⎨ .

⎭ ⎩Na 3 VO 4

In the case of others, it is necessary to have oxidizing agent:

Cr + NaNO 3 + NaOH→ Na 2 CrO 4 + NaNO 2,

O 2 + Na 2 CO 3 → Na 2 WO 4 + CO 2 .

Moreover, W and Mo interact with alkali more actively than Cr, because During the reaction, their surface is covered with a more acidic oxide (EO) 3 than in the case of chromium (Cr 2 O 3).

Interaction of d-metals with simple substances. Corrosion. At room conditions, only fluorine oxidizes most d-metals, except noble ones (but reactions with Cu, Ni, Fe (as well as with Pb, Al) are limited to the formation of protective films of fluorides). In addition, at ob.u. gold interacts with bromine, and mercury interacts with iodine and sulfur due to the formation of thermodynamically very sustainable products: AuBr, 3 HgI 2 and HgS (see section “Halogens”).

In air, in a finely dispersed state, fairly active metals (Ti, Cr, Mn, Fe, Co, Ni) pyrophores 2 (i.e. they light up when exposed to air), but in compact form most M are stable due to passivation. Especially dense surface films are formed by metals of the vanadium and titanium subgroup, therefore they have high corrosion resistance (even in sea water).

Other metals are not as stable. Under the influence of air components (which ones?), corrosion of zinc and copper occurs slowly (with the formation of E 2 (OH) 2 CO 3); Even silver darkens, becoming covered with sulfide (under the combined influence of O 2, H 2 O and H 2 S; what is the role of each of them?).

Iron corrodes especially quickly. True, in a dry atmosphere its oxidation occurs only before the formation dense FeO passivating film. But in the presence of moisture, the product obtained by the reaction:

Fe + H 2 O → FeO+ H 2 ,

oxidized by oxygen, activated by H 2 O molecules, to Fe 2 O 3. In this case, water sorbed by the metal surface, partially dissolving oxidation products in itself, hinders formation dense oxide structure, as a result of which iron corrosion occurs deep down.

The addition of alkali reduces the oxidation potential of oxygen, and therefore the process proceeds to a lesser extent. Note that Veryclean iron, which adsorbs hydrogen well and thus passivates its surface, does not oxidize.

To protect against corrosion, industrial iron is painted or subjected to tinning, galvanizing, chrome plating, nickel plating, nitriding (Fe 4 N coating), cementing (Fe C 3) and other processing methods. In particular, vitrification metal surface treatment with a laser increases corrosion resistance 12 times, but when M is heated above 200 0 C, this effect is lost. A more reliable, but expensive way to combat iron oxidation in air is to produce stainless steel (18% Cr and 9% Ni).

However, corrosion is a slow process, and quite fast d-metals react with nonmetals only when heated, even the most active M subgroups of scandium (oxidizing to (+3)). (However, from Sc to La the interaction activity increases (?), and lanthanum, for example, ignites in chlorine at ob.u.)

In the case of less reactive (?) metals of the titanium subgroup, it is required more heating (above 150 0 C). In this case, Hf transforms into Hf + 4, and Ti and Zr can form products in inferior st.ok.: Ti 2 O 3, ZrCl 2, etc. However, they are strong reducing agents, especially in the case of Zr (?) - they oxidize in air or dismutate:

ZrCl 2 → Zr+ ZrCl 4 .

With even less active metals of the vanadium subgroup, reactions occur at t > 400 0 C, and with the formation of products only at the highest degree. (+5).

When moving to the chromium subgroup, the reactivity M growing(due to the greater volatility of the oxides), but decreases from Cr to W (?). Thus, chromium interacts with all Г2, molybdenum does not react with I2, and tungsten does not react with Br2. Moreover, the oxidation of chromium goes up to (+3), and its analogues - up to (+6). (Note that WF is the 6th heaviest gas at zero level)

Similar patterns are observed in other subgroups of d-metals. Thus, technetium and rhenium do not interact with iodine, and with other halogens - only at t > 400 0 C, forming EG 7. At the same time, manganese oxidizes with slight heating

even gray and up to st.ok. (+2).

Copper reacts with wet chlorine at r.p.c., silver - with slight heating, and gold - only at t> 200 0 C. When heated, oxygen acts only on copper (CuO product, at higher temperatures - Cu 2 O (?)), and silver oxidized (unlike gold) by ozone (to AgO).

Zinc also burns in CO 2 , and mercury at ambient conditions. It is not even covered with an oxide film. When heated to 300 0 C, it forms a mixture of oxides HgO and Hg 2 O, which at t> 400 0 C split off O, turning into Hg, while the decomposition temperature of cadmium oxide is 1813 0 C, and ZnO is 1950 0 C.

The most chemically stable platinum metals and gold, but with sufficient heating they react with almost all non-metals (G 2, O 2, S, P, As), although with different activity and selectivity; namely: in periods from left to right, resistance to O 2 and F 2 increases, and to Cl 2 and S decreases (in accordance with the electronic structure of the atoms of the elements (?)).

So, if fluorine reacts with platinum only at t > 400 0 C, then chlorine reacts at 250 0 C (product PtCl 2). Or if we consider the interaction with oxygen: osmium in the form of black is oxidized in air at r.b. (up to OsO 4), ruthenium - with slight heating, and the rest - at a red heat temperature. Products: IrO 2, PdO, PtO 2, Rh 2 O 3.

(With stronger heating, these oxides decompose, and if the reaction:

PtO 2 → Pt+ O 2

occurs at 500 0 C, then decomposition:

RuO 2 → Ru+ O 2

occurs only when t > 1300 0 C).

A similar increase in the metal’s resistance to oxygen is observed when moving from iron to nickel (see Table 14).

Table 14. Characteristics of the interaction of metals of the iron family with oxygen

Formation of solid solutions. A feature of d-metals is their tendency due to the wide variety of st.ok. and valence states to form compounds non-stoichiometric composition: intermetallic compounds (AlNi, etc.) or metallides (Fe S 3, VN, LaB, ZrC 6, etc.). And solid solutions, in particular, solutions implementation gases Thus, metals of the scandium and titanium subgroup absorb hydrogen at r.p.a. to the composition: EH 2 and EH (3 when heated, the solubility of H 2 decreases).

Nickel and palladium have a special affinity for hydrogen (1 V Pd dissolves 1000 V H 2), which are therefore reaction catalysts hydrogenation. And, for example, platinum predominantly sorbs O2 (up to 700 V) and is therefore used as a catalyst for processes involving oxygen: oxidation NH 3 to NO, SO 2 to SO, 3 for afterburning car exhaust gases (in this case, in particular, NO turns into N 2, and CO into CO 2), etc.

The mechanism of the catalytic action of these metals is that, as is assumed, gases dissolving in M atomized. Thus, hydrogen released when its solution in a metal is heated is a stronger reducing agent than molecular.

In addition, for example, palladium, when absorbing H 2 to a certain limit, retains its metallic properties, but loses paramagnetism. This means that at least some of the hydrogen atoms give up their valence electrons to the conduction band of the metal.

There is also evidence of partial formation of hydride ions, for example, when hydrogen dissolves in iron. Received, etc. unconventional hydrides in which H 2 molecules are coordinated as a whole on the d-metal atom. (They serve as models for studying intermediates that arise during catalysis.)