Alkali metals- these are the elements of the 1st group of the periodic table of chemical elements (according to the outdated classification - elements of the main subgroup of group I): lithium Li, sodium Na, potassium K, rubidium Rb, cesium Cs, France Fr, and despondent Uue. When alkali metals dissolve in water, soluble hydroxides are formed, called alkalis.
Chemical properties of alkali metals
Due to the high chemical activity of alkali metals towards water, oxygen, and sometimes even nitrogen (Li, Cs), they are stored under a layer of kerosene. To carry out a reaction with an alkali metal, a piece of the required size is carefully cut off with a scalpel under a layer of kerosene, the metal surface is thoroughly cleaned in an argon atmosphere from the products of its interaction with air, and only then the sample is placed in the reaction vessel.
1. Interaction with water. An important property of alkali metals is their high activity towards water. Lithium reacts most calmly (without explosion) with water:
When a similar reaction is carried out, sodium burns with a yellow flame and a small explosion occurs. Potassium is even more active: in this case, the explosion is much stronger, and the flame is colored purple.
2. Interaction with oxygen. The combustion products of alkali metals in air have different compositions depending on the activity of the metal.
· Only lithium burns in air to form an oxide of stoichiometric composition:
· When burning sodium mainly Na 2 O 2 peroxide is formed with a small admixture of NaO 2 superoxide:
· In combustion products potassium, rubidium And cesium contains mainly superoxides:
To obtain sodium and potassium oxides, mixtures of hydroxide, peroxide or superoxide with an excess of metal are heated in the absence of oxygen:
The following pattern is characteristic of oxygen compounds of alkali metals: as the radius of the alkali metal cation increases, the stability of oxygen compounds containing peroxide ion O 2 2− and superoxide ion O 2− increases.
Heavy alkali metals are characterized by the formation of fairly stable ozonides composition of EO 3. All oxygen compounds have different colors, the intensity of which deepens in the series from Li to Cs:
Alkali metal oxides have all the properties of basic oxides: they react with water, acidic oxides and acids:
Peroxides And superoxides exhibit the properties of strong oxidizing agents:
Peroxides and superoxides interact intensively with water, forming hydroxides:
3. Interaction with other substances. Alkali metals react with many nonmetals. When heated, they combine with hydrogen to form hydrides, with halogens, sulfur, nitrogen, phosphorus, carbon and silicon to form, respectively, halides, sulfides, nitrides, phosphides, carbides And silicides:
When heated, alkali metals are capable of reacting with other metals, forming intermetallic compounds. Alkali metals react actively (explosively) with acids.
Alkali metals dissolve in liquid ammonia and its derivatives - amines and amides:
When dissolved in liquid ammonia, an alkali metal loses an electron, which is solvated by ammonia molecules and gives the solution a blue color. The resulting amides are easily decomposed by water to form alkali and ammonia:
Alkali metals interact with organic substances, alcohols (to form alcoholates) and carboxylic acids (to form salts):
4. Qualitative determination of alkali metals. Since the ionization potentials of alkali metals are small, when the metal or its compounds are heated in a flame, the atom is ionized, coloring the flame a certain color:
Flame coloring with alkali metals
and their connections
Alkaline earth metals.
Alkaline earth metals- chemical elements of group II of the periodic table of elements: beryllium, magnesium, calcium, strontium, barium and radium.
Physical properties
All alkaline earth metals are gray substances that are solid at room temperature. Unlike alkali metals, they are significantly harder and generally cannot be cut with a knife (with the exception of strontium). The density of alkaline earth metals with atomic number increases, although growth is clearly observed only starting with calcium, which has the lowest density among them (ρ = 1.55 g/cm³), the heaviest is radium, the density of which is approximately equal to the density of iron.
Chemical properties
Alkaline earth metals have an outer energy level electronic configuration ns², and are s-elements, along with the alkali metals. Having two valence electrons, alkaline earth metals easily give them up, and in all compounds they have an oxidation state of +2 (very rarely +1).
The chemical activity of alkaline earth metals increases with increasing atomic number. Beryllium in its compact form does not react with oxygen or halogens, even at red heat temperatures (up to 600 °C; reactions with oxygen and other chalcogens require an even higher temperature, fluorine is an exception). Magnesium is protected by an oxide film at room temperature and higher temperatures (up to 650 °C) and does not oxidize further. Calcium oxidizes slowly and deeply at room temperature (in the presence of water vapor), and burns with slight heating in oxygen, but is stable in dry air at room temperature. Strontium, barium and radium quickly oxidize in air, giving a mixture of oxides and nitrides, so they, like alkali metals (and calcium), are stored under a layer of kerosene.
Oxides and hydroxides of alkaline earth metals tend to increase their basic properties with increasing atomic number: Be(OH) 2 is an amphoteric, water-insoluble hydroxide, but soluble in acids (and also exhibits acidic properties in the presence of strong alkalis), Mg(OH) 2 - weak base, insoluble in water, Ca(OH) 2 - strong but slightly soluble base in water, Sr(OH) 2 - more soluble in water than calcium hydroxide, strong base (alkali) at high temperatures close to the boiling point water (100 °C), Ba(OH) 2 is a strong base (alkali), not inferior in strength to KOH or NaOH, and Ra(OH) 2 is one of the strongest alkalis, a very corrosive substance
Being in nature
All alkaline earth metals are found (in varying quantities) in nature. Due to their high chemical activity, all of them are not found in a free state. The most common alkaline earth metal is calcium, the amount of which is 3.38% (by weight of the earth’s crust). It is slightly inferior to magnesium, the amount of which is 2.35% (of the mass of the earth’s crust). Barium and strontium are also common in nature, accounting for 0.05 and 0.034% of the mass of the earth's crust, respectively. Beryllium is a rare element, the amount of which is 6·10−4% of the mass of the earth's crust. As for radium, which is radioactive, it is the rarest of all alkaline earth metals, but it is always found in small quantities in uranium ores. In particular, it can be isolated from there chemically. Its content is 1·10−10% (of the mass of the earth’s crust)
Aluminum.
Aluminum- an element of the main subgroup of the third group of the third period of the periodic system of chemical elements of D. I. Mendeleev, with atomic number 13. Denoted by the symbol Al(lat. Aluminum). Belongs to the group of light metals. The most common metal and the third most abundant chemical element in the earth's crust (after oxygen and silicon).
Simple substance aluminum- a lightweight, paramagnetic metal of silver-white color, easy to form, cast, and machine. Aluminum has high thermal and electrical conductivity and resistance to corrosion due to the rapid formation of strong oxide films that protect the surface from further interaction.
Aluminum was first obtained by the Danish physicist Hans Oersted in 1825 by the action of potassium amalgam on aluminum chloride followed by distillation of mercury. The modern production method was developed independently by the American Charles Hall and the Frenchman Paul Héroult in 1886. It consists of dissolving aluminum oxide Al 2 O 3 in a melt of cryolite Na 3 AlF 6 followed by electrolysis using consumable coke or graphite electrodes. This production method requires a lot of electricity, and therefore became popular only in the 20th century.
To produce 1000 kg of crude aluminum, 1920 kg of alumina, 65 kg of cryolite, 35 kg of aluminum fluoride, 600 kg of anode mass and 17 thousand kWh of DC electricity are required
Special (correctional)
comprehensive boarding school for the blind
and visually impaired children in Perm
Abstract completed
10th grade students
Ponomarev Oleg,
Korshunov Artem
Supervisor:
L.Yu. Zakharova,
chemistry teacher
Perm
Introduction
General characteristics of elements of group I A-group
4 – 10
1.1. History of the discovery and distribution of alkali metals in nature
4 – 5
5 - 6
6 – 8
8 – 9
9 – 10
Biological role of elements of group I A-group. Their use in medicine
11 – 17
Routes of entry of alkali metals into the human body
18 – 21
Practical work
22 – 23
conclusions
24 – 25
Used Books
Introduction
The time has long come when everyone should think about their health and not only their own. We do not very often use the knowledge we acquire at school, for example in chemistry, in everyday life. However, this particular subject can become a source of knowledge about our health. Thanks to chemistry, we learn how the substances of our planet affect the vital processes of the body, and in general human life itself, what is useful to us and in what quantities and, finally, what is harmful and to what extent.
The human body is a complex chemical system that cannot function independently, without connection with the environment. It has been proven that almost all chemical elements are present in a living organism: some are macroelements, while the content of others is negligible, these are microelements. The ways in which elements enter the body are different, and their influence on the body is varied, but each plays its own biological role.
It is impossible to study the meaning of each element within the framework of one work. We have chosen the very first group of chemical elements of D.I. Mendeleev’s periodic table.
Target of this study – study the biological role of alkali metals for the human body.
In this regard, we decided to clarify the following questions for each metal of group IA:
general characteristics and structural features of the atoms of each element, as well as the properties of the substances they form;
presence of the element in the body;
the body's needs for it;
the effect of excess and deficiency of the element on human health;
natural sources;
methods for detecting an element.
1. General characteristics of elements of group I A-group
Period
Group
IN I A-group includes s-elements - alkali metals, which are extremely important for the normal life of animals and people. The macroelements sodium and potassium are of greatest importance for organisms.
3Li
11 Na
19K
37 Rb
55 Cs
87 Fr
1.1. History of discovery and distribution in nature
alkali metals
The name “alkali metals” is due to the fact that the hydroxides of the two main representatives of this group - sodium and potassium - have long been known as alkalis. From these alkalis, subjecting them to electrolysis in a molten state, G. Davy in 1807 for the first time received free potassium and sodium. J. Berzelius proposed to name element No. 11 sodium (from Arabic natrun- soda), and element No. 19, at Gilbert’s suggestion, was named potassium (from the Arabic alkali– alkali).
The remaining metals were isolated by scientists from the compounds later. Lithium was discovered by the Swedish chemist I. Arfvedson in 1817, and at the suggestion of J. Berzelius it was called lithium (from the Greek litos- stone), because Unlike potassium, which until then had only been found in plant ashes, it was found in stone.
Rubidium was isolated in 1861, cesium in 1860. Francium was obtained artificially in 1939. French researcher M. Pere during the decay of actinium, is a radioactive element.
Due to their very easy oxidation, alkali metals occur in nature exclusively in the form of compounds. Some of their natural compounds, in particular sodium and potassium salts, are quite widespread; they are found in many minerals, plants, and natural waters.
Sodium and potassium are common elements: the content of each of them in the earth's crust is approximately 2% by weight. Both metals are found in various minerals and silicate-type burrow rocks.
Sodium chloride NaCl is found in seawater and also forms thick rock salt deposits in many places around the globe. The upper layers of these deposits sometimes contain quite significant amounts of potassium, mainly in the form of chloride KCl or double salts with sodium and magnesium KCl ∙MgCl 2. However, large accumulations of potassium salts of industrial importance are rare. The most important of them are the Solikamsk deposits (sylvinite) in Russia, the Strassfurt deposits in Germany and the Alsatian deposits in France.
Deposits of sodium nitrate NaNO 3 are located in Chile. The water of many lakes contains Na 2 CO 3 soda. Finally, huge quantities of sodium sulfate Na 2 SO 4 are found in the Kara-Bogaz-Gol Bay of the Caspian Sea, where this salt is deposited in a thick layer on the bottom during the winter months.
Lithium, rubidium and cesium are much less common than sodium and potassium. Lithium is the most common, but minerals containing it rarely form large accumulations. Rubidium and cesium are found in small quantities in some lithium minerals.
Francium is found in nature in insignificant quantities (there is hardly 500g of it on the entire globe); it is obtained artificially.
1.2. Structure and properties of alkali metal atoms
The electronic formula of the valence shell of alkali metal atoms is ns 1, i.e. the atoms of these elements have one valence electron in the s sublevel of the outer energy level. Accordingly, the stable oxidation state of alkali metals is +1.
All elements of the IA group are very similar in properties, which is explained by the similar structure of not only the valence electron shell, but also the outer one (with the exception of lithium).
As the radius of an atom in the Li – Na – K – Rb – Cs – Fr group increases, the bond between the valence electron and the nucleus weakens. Accordingly, in this series the ionization energy of alkali metal atoms decreases.
Having one electron in their valence shells, located at a great distance from the nucleus, alkali metal atoms easily give up an electron. This causes low ionization energy. As a result of ionization, E + cations are formed, which have a stable electronic configuration of noble gas atoms.
The table shows some properties of alkali metal atoms.
Characteristic
3 Li
11 Na
1 9K
37 Rb
55 Cs
87 Fr
Valence electrons
2s 1
3s 1
4s 1
5s 1
6s 1
7s 1
Molar mass, g/mol
23,0
39,1
85,5
132,9
Metallic radius of an atom, pm
Crystal radius of an atom, pm
Ionization energy,
kJ/mol
Alkali metals are the most typical representatives of metals: their metallic properties are especially pronounced.
1.3. Alkali metals are simple substances
Silvery-white soft substances (cut with a knife), with a characteristic shine on the freshly cut surface. When exposed to air, the shiny surface of the metal immediately becomes dull due to oxidation.
All of them are light and fusible, and, as a rule, their density increases from Li to Cs, and the melting point, on the contrary, decreases.
Characteristic
Li
Na
K
Rb
Cs
Fr
Density, g/cm 3
0,53
0,97
0,86
1,53
Hardness (diamond = 10)
Electrical conductivity (Hg = 1)
11,2
13,6
Melting point, C
Boiling point, C
1350
Standard electrode potential, V
3,05
2,71
2,92
2,93
2,92
Coordination number
4, 6
4, 6
6, 8
All alkali metals have negative standard redox potentials, large in absolute value. This characterizes them as very strong reducing agents. Only lithium is somewhat inferior to many metals in chemical activity.
Despite the similarity of properties, sodium and especially lithium differ from other alkali metals. The latter is primarily due to the significant difference in the radii of their atoms and the structure of the electron shells.
Alkali metals are among the most chemically active elements. The chemical activity of alkali metals naturally increases with increasing atomic radius.
Li Na K Rb Cs Fr
Chemical activity increases
The radius of the atom increases
Alkali metals actively interact with almost all non-metals.
When interacting with oxygen lithium forms the oxide Li 2 O, and the remaining alkali metals form peroxides Na 2 O 2 and superoxides KO 2, RbO 2, CsO 2. For example:
4Li (t) + O 2 (g) = 2Li 2 O (t)
2Na (t) + O 2 (g) = Na 2 O 2 (t)
K (t) + O 2 (g) = KO 2 (t)
Alkali metals react actively with halogens, forming EG halides; with sulfur- with the formation of E 2 S sulfides. Alkali metals, with the exception of lithium, do not react directly with nitrogen.
2E(t) + Cl 2 (g) = 2ECl (t)
2E(t) + S (t) = E 2 S (t)
All alkali metals react directly with water, forming EON hydroxides - alkalis and reducing water to hydrogen:
2E (t) + 2H 2 O (l) = 2EON (r) + H 2 (g)
The intensity of interaction with water increases significantly in the Li-Cs series.
The reducing power of alkali metals is so great that they can even reduce hydrogen atoms, turning them into negatively charged H - ions. Thus, when heating alkali metals in a jet hydrogen their hydrides are obtained, for example:
2E(t) + N 2 (g) = 2EN
1.4. Application of alkali metals
Alkali metals and their compounds are widely used in technology.
Lithium is used in nuclear energy. In particular, the 6 Li isotope serves as an industrial source for the production of tritium, and the 7 Li isotope is used as a coolant in uranium reactors. Due to the ability of lithium to easily combine with hydrogen, nitrogen, oxygen, and sulfur, it is used in metallurgy to remove traces of these elements from metals and alloys.
Lithium and its compounds are also used as fuel for rockets. Lubricants containing lithium compounds retain their properties over a wide temperature range. Lithium is used in ceramics, glass and other chemical industries. In general, in terms of importance in modern technology, this metal is one of the most important rare elements.
Cesium and rubidium are used to make solar cells. These devices, which convert radiant energy into electric current energy and are based on the phenomenon of the photoelectric effect, use the ability of cesium and rubidium atoms to split off valence electrons when exposed to radiant energy on the metal.
The most important areas of application of sodium are nuclear energy, metallurgy, and the organic synthesis industry.
In nuclear energy, sodium and its alloy with potassium are used as liquid metal coolants. The sodium-potassium alloy, containing 77.2% potassium, is in a liquid state over a wide temperature range, has a high heat transfer coefficient and does not interact with most structural materials.
In metallurgy, a number of refractory metals are obtained by the sodium thermal method. In addition, sodium is used as an additive to strengthen lead alloys.
In the organic synthesis industry, sodium is used in the production of many substances. It also serves as a catalyst in the production of some organic polymers.
Potassium is one of the elements required in significant quantities for plant nutrition. Although there is quite a lot of potassium salts in the soil, a lot of it is also carried away with some cultivated plants. Flax, hemp and tobacco carry away especially much potassium. To replenish the loss of potassium from the soil, it is necessary to add potassium fertilizers to the soil.
1.5. Alkali metal compounds
Oxides E 2 ABOUT- solids. They have pronounced basic properties: they interact with water, acids and acid oxides. For example:
E 2 O(t) + H 2 O(l) = 2EON (p)
Peroxides and superoxides E 2 ABOUT 2 and EO 2 alkali metals are strong oxidizing agents. Sodium peroxide and potassium superoxide are used in closed objects (submarines, spacecraft) to absorb carbon dioxide and regenerate oxygen:
2Na 2 O 2 (t) + 2CO 2 (g) = 2Na 2 CO 3 (t) + O 2 (g)
4KO 2 (t) + 2CO 2 (g) = 2K 2 CO 3 (t) + 3O 2 (g)
Sodium peroxide is also used to bleach fabrics, wool, silk, etc.
Alkalis– solid, white, very hygroscopic crystalline substances, relatively fusible and highly soluble in water (with the exception of LiOH). Solid alkalis and their concentrated solutions have a corrosive effect on fabrics, paper and living tissues due to dehydration and alkaline hydrolysis of proteins. Therefore, working with them requires protective precautions. Due to their strong corrosive effect, these alkalis are called caustic (NaOH - caustic soda, caustic, KOH - caustic potassium).
Alkalies dissolve well in water with the release of a large amount of heat, exhibit pronounced properties of strong soluble bases: they interact with acids, acid oxides, salts, amphoteric oxides and hydroxides.
Caustic soda is used in large quantities to purify petroleum products. in the paper and textile industries, for the production of soap and fibers.
Caustic potassium is more expensive and is used less frequently. Its main area of application is the production of liquid soap.
Alkali metal salts– solid crystalline substances of ionic structure. The most important of them are carbonates, sulfates, and chlorides.
Most alkali metal salts are highly soluble in water (with the exception of lithium salts: Li 2 CO 3, LiF, Li 3 PO 4).
With polybasic acids, alkali metals form both medium (E 2 SO 4, E 3 PO 4, E 2 CO 3, E 2 SO 3, etc.) and acidic (ENSO 4, EN 2 PO 4, E 2 NPO 4, ENSO 3, etc.) salts.
Na 2 CO 3 - sodium carbonate, forms crystalline hydrate Na 2 CO 3 ∙10H 2 CO 3, known as crystalline soda, which is used in the production of glass, paper, and soap. This is medium salt.
In everyday life, the more commonly known acidic salt is sodium bicarbonate NaHCO 3; it is used in the food industry (baking soda) and in medicine (baking soda).
K 2 CO 3 - potassium carbonate, technical name - potash, used in the production of liquid soap and for the preparation of refractory glass, and also as a fertilizer.
Na 2 SO 4 ∙10H 2 O – sodium sulfate crystalline hydrate, technical name Glauber’s salt, is used for the production of soda and glass, and also as a laxative.
NaCl - sodium chloride, or table salt, is the most important raw material in the chemical industry and is widely used in everyday life.
2. Biological role of s-elements of group IA. Their use in medicine
Chemical element, E
10 -4 %
0,08%
0,23%
10 -5 %
10 -4 %
Alkali metals in the form of various compounds are part of human and animal tissues.
Sodium and potassium are vital elements that are constantly present in the body and participate in metabolism. Lithium, rubidium and cesium are also constantly contained in the body, but their physiological and biochemical role is poorly understood. They can be classified as trace elements.
In the human body, alkali metals are found in the form of the E + cation.
The similarity of the electronic structure of alkali metal ions, and, consequently, the physicochemical properties of the compounds also determines the similarity of their effect on biological processes. Differences in electronic structure determine their different biological roles. On this basis, it is possible to predict the behavior of alkali metals in living organisms.
Thus, sodium and lithium accumulate in the extracellular fluid, and potassium, rubidium and cesium accumulate in the intracellular fluid. Lithium and sodium are especially close in biological action. For example, they are very similar in their enzyme-activating properties.
The similarity of the properties of sodium and lithium determines their interchangeability in the body. In this regard, with excessive introduction of sodium or lithium ions into the body, they are able to equivalently replace each other. This is the basis for the administration of sodium chloride in cases of lithium salt poisoning. In accordance with Le Chatelier's principle, the balance between sodium and lithium ions in the body shifts towards the elimination of Li + ions, which leads to a decrease in its concentration and the achievement of a therapeutic effect.
Rubidium and cesium are close in physical and chemical properties to potassium ions, so they behave in a similar way in living organisms. In the systems studied, potassium, rubidium and cesium are synergists, and with lithium they are antagonists. The similarity of rubidium and potassium is the basis for the introduction of potassium salts into the body in case of poisoning with rubidium salts.
Sodium and potassium, as a rule, are antagonists, but in some cases the similarity of many physicochemical properties determines their interchange in living organisms. For example, with an increase in the amount of sodium in the body, the excretion of potassium by the kidneys increases, i.e., hypokalemia occurs.
Lithium. The lithium content in the human body is about 70 mg (10 mmol). Lithium is one of the most valuable microelements, or, as they also call it, mini-metals. Lithium was once used to treat gout and eczema. And in 1971 An interesting message appeared in the magazine “Medical News”: in those areas where drinking water contains large amounts of lithium, people are kinder and calmer, there are fewer rude people and brawlers among them, and there are significantly fewer mental illnesses. The psychotropic properties of this metal were revealed. Lithium began to be used for depression, hypochondria, aggressiveness and even drug addiction.
However, lithium can be both “good” and “evil”. There have been cases when, during injection treatment with lithium, a powerful metabolic disorder occurred, and serious consequences of this are inevitable.
Lithium compounds in higher animals are concentrated in the liver, kidneys, spleen, lungs, blood, and milk. The maximum amount of lithium is found in human muscles. The biological role of lithium as a trace element has not yet been fully elucidated.
It has been proven that at the level of cell membranes, lithium ions compete with sodium ions when entering cells. Obviously, the replacement of sodium ions in cells with lithium ions is associated with a greater covalency of lithium compounds, as a result of which they are better soluble in phospholipids.
It has been established that some lithium compounds have a positive effect on patients with manic depression. Absorbed from the gastrointestinal tract, lithium ions accumulate in the blood. When the concentration of lithium ions reaches 0.6 mmol/l and above, there is a decrease in emotional tension and a weakening of manic excitement. However, the content of lithium ions in the blood plasma must be strictly controlled. In cases where the concentration of lithium ions exceeds 1.6 mmol/l, negative phenomena are possible.
It is now known that in addition to psychotropic effects, lithium has properties to prevent sclerosis, heart disease, and to some extent diabetes and hypertension. It “helps” magnesium in its anti-sclerotic protection.
At the end of 1977 The results of studies conducted at the Krakow Hematology Clinic were published. The studies were devoted to the influence of lithium on the hematopoietic system. It turned out that this microelement activates the action of bone marrow cells that have not yet died. The discovery could play an important role in the fight against blood cancer. Research is still ongoing. I would like to believe that their results will bring invaluable help to people.
Sodium. The sodium content in the human body weighing 70 kg is about 60 g (2610 mmol). Of this amount, 44% of sodium is in the extracellular fluid and 9% in the intracellular fluid.
The remaining amount of sodium is found in bone tissue, which is the site of deposition of Na + ion in the body. About 40% of the sodium contained in bone tissue is involved in metabolic processes and due to this, the skeleton is either a donor or acceptor of sodium ions, which helps maintain a constant concentration of sodium ions in the extracellular fluid.
Sodium is the main extracellular ion. The human body contains sodium in the form of its soluble salts, mainly NaCl chloride, Na 3 PO 4 phosphate and NaHCO 3 bicarbonate.
Sodium is distributed throughout the body: in blood serum, cerebrospinal fluid, eye fluid, digestive juices, bile, kidneys, skin, bone tissue, lungs, brain.
Sodium ions play an important role in ensuring the constancy of the internal environment of the human body, participates in maintaining a constant osmotic pressure of the biofluid, and ensures the acid-base balance of the body. Sodium ions are involved in the regulation of ion exchange and affect the functioning of enzymes. Together with potassium, magnesium, calcium, and chlorine ions, sodium ion participates in the transmission of nerve impulses through the membranes of nerve cells and maintains normal excitability of muscle cells.
When the sodium content in the body changes, dysfunctions of the nervous, cardiovascular and other systems, smooth and skeletal muscles occur. Sodium chloride NaCl serves as the main source of hydrochloric acid for gastric juice.
Sodium enters the human body mainly in the form of table salt NaCl. The body's true daily need for sodium is 1g, although the average consumption of this element reaches 4 - 7g.
Continuous excess consumption of NaCI contributes to the appearance of hypertension. In the body of a healthy person, a balance is maintained between the amount of sodium consumed and excreted. About 90% of sodium consumed is excreted in urine, and the rest in sweat and feces.
So, to summarize: sodium ions play an important role:
to ensure osmotic homeostasis
to ensure the acid-base balance of the body
in the regulation of water metabolism
in the work of enzymes
in the transmission of nerve impulses
in the work of muscle cells
Isotonic solutionNaCI (0,9%) for injection, it is administered subcutaneously, intravenously and in enemas for dehydration and intoxication, and is also used for washing wounds, eyes, nasal mucosa, as well as for dissolving various medications.
Hypertonic solutionsNaCI (3-5-10%) used externally in the form of compresses and lotions in the treatment of purulent wounds. The use of such compresses promotes, by the law of osmosis, the separation of pus from wounds and plasmolysis of bacteria (antimicrobial effect). A 2-5% NaCI solution is prescribed orally for gastric lavage in case of AgNO 3 poisoning, which is converted into slightly soluble and non-toxic silver chloride:
Ag + + CI - = AgCI (t)
Drinking soda(sodium bicarbonate, bicarbonate of soda) NaHCO 3 is used for various diseases accompanied by high acidity - acidosis (diabetes, etc.). The mechanism for reducing acidity is the interaction of NaHCO 3 with acidic products. In this case, sodium salts of organic acids are formed, which are largely excreted in the urine, and carbon dioxide, which leaves the body with exhaled air:
NaHCO3 (p) + RCOOH (p) → RCOONa(p) + H 2 O(l) + CO 2 (g)
NaHCO 3 is also used for increased acidity of gastric juice, gastric and duodenal ulcers. When taking NaHCO 3, a neutralization reaction of excess hydrochloric acid occurs:
NaHCO 3 (s) + HCl (s) = NaCl (s) + H 2 O (l) + CO 2 (g)
It should be borne in mind that the use of baking soda should be careful, because... may cause a number of side effects.
Solutions of baking soda are used as rinses and washes for inflammatory diseases of the eyes and mucous membranes of the upper respiratory tract. The action of NaHCO 3 as an antiseptic is based on the fact that, as a result of hydrolysis, an aqueous soda solution exhibits slightly alkaline properties:
NaHCO 3 + H 2 O ↔ NaOH + H 2 CO 3
When microbial cells are exposed to alkalis, precipitation of cellular proteins occurs and, as a result, the death of microorganisms.
Glauber's salt(sodium sulfate) Na 2 SO 4 ∙10H 2 O is used as a laxative. This salt is slowly absorbed from the intestine, which leads to the maintenance of increased osmotic pressure in the intestinal cavity for a long time. As a result of osmosis, water accumulates in the intestines, its contents become liquefied, intestinal contractions intensify, and feces are eliminated faster.
Borax(sodium tetraborate) Na 2 B 4 O 7 ∙10H 2 O is used externally as an antiseptic for rinsing, douching, and lubricating. the antiseptic effect of borax is similar to the effect of baking soda and is associated with the alkaline reaction of the aqueous solution of this salt, as well as with the formation of boric acid:
Na 2 B 4 O 7 + 7H 2 O ↔ 4H 3 BO 3 + 2NaOH
Sodium hydroxide in the form of a 10% NaOH solution, it is included in the composition of silane, used in orthopedic practice for casting fire-resistant models in the manufacture of solid prostheses from a cobalt-chrome alloy.
Radioactive isotope 24 Na is used as a tracer to determine the speed of blood flow, and it is also used to treat some forms of leukemia.
Potassium. The potassium content in the human body weighing 70 kg is approximately 160 g (4090 mmol). Potassium is the main intracellular cation, accounting for 2/3 of the total active cellular cations. In most cases, potassium is an antagonist to sodium.
Of the total amount of potassium contained in the body, 98% is found inside cells and only about 2% is in extracellular fluid. Potassium is distributed throughout the body. Its topography: liver, kidneys, heart, bone tissue, muscles, blood, brain, etc.
Potassium ions K+ play an important role in physiological processes:
muscle contraction
in the normal functioning of the heart
participates in the transmission of nerve impulses
in exchange reactions
activates the work of a number of enzymes located inside the cell
regulates acid-base balance
It has protective properties against the unwanted effects of excess sodium and normalizes blood pressure. In the body of people who eat a lot of potassium-rich vegetables - vegetarians - the amount of potassium and sodium are in balance. These people most often have lower blood pressure than their meat-loving fellow citizens.
Has an antisclerotic effect
Potassium has the ability to enhance urine formation
An adult usually consumes 2–3 g of potassium per day with food. The concentration of potassium ions in extracellular fluid, including plasma, is normally 3.5 - 5.5 mmol/l, and the concentration of intracellular potassium is 115 - 125 mmol/l.
Rubidium and cesium. According to their content in the human body, rubidium and cesium are classified as microelements. They are constantly contained in the body, but their biological role has not yet been clarified.
Rubidium and cesium are found in all studied organs of mammals and humans. Entering the body with food, they are quickly absorbed from the gastrointestinal tract into the blood. The average level of rubidium in the blood is 2.3-2.7 mg/l, and its concentration in erythrocytes is almost three times higher than in plasma. Rubidium and cesium are distributed very evenly in organs and tissues, and rubidium mainly accumulates in the muscles, and cesium enters the intestine and is reabsorbed in its descending sections.
The role of rubidium and cesium in some physiological processes is known. Currently, the stimulating effect of these elements on circulatory functions and the effectiveness of the use of their salts for hypotension of various origins have been established. In the laboratory of I.P. Pavlov, S.S. Botkin found that cesium and rubidium chlorides cause an increase in blood pressure for a long time and that this effect is associated mainly with increased cardiovascular activity and constriction of peripheral vessels.
Being a complete analogue of potassium, rubidium also accumulates in intracellular fluid and can replace an equivalent amount of potassium in various processes. Synergism (chemical) is the simultaneous combined effect of two (or more) factors, characterized by the fact that such a combined effect significantly exceeds the effect of each individual component. A potassium synergist, rubidium activates many of the same enzymes as potassium.
Radioactive isotopes 137 Cs and 87 Rb are used in radiotherapy of malignant tumors, as well as in the study of potassium metabolism. Due to their rapid breakdown, they can even be introduced into the body without fear of long-term harmful effects.
Franc. It is a radioactive chemical element obtained artificially. There is evidence that francium is capable of selectively accumulating in tumors at the earliest stages of their development. These observations may be useful in diagnosing cancer.
Thus, Of the IA group elements, Li, Rb, Cs are physiologically active, and Na and K are vital. The similarity of the physicochemical properties of Li and Na, due to the similarity of the electronic structure of their atoms, is also manifested in the biological action of cations (accumulation in the extracellular fluid, interchangeability). The similar nature of the biological action of cations of elements of long periods - K +, Rb +, Cs + (accumulation in intracellular fluid, interchangeability) is also due to the similarity of their electronic structure and physicochemical properties. This is the basis for the use of sodium and potassium preparations for poisoning with lithium and rubidium salts.
3. Routes of entry of alkali metals
into the human body
The ways in which chemical elements enter the human body are varied; they are presented in the diagram:
Human
In the process of evolution from inorganic to bioorganic substances, the basis for the use of certain chemical elements in the creation of biological systems is natural selection.
The table shows data on the content of group I A elements - alkali metals - in the earth's crust, sea water, plant and animal organisms and in the human body (mass fraction in %).
The table shows that the greater the abundance of an element in the earth’s crust, the more it is in the human body.
Li
Na
K
Rb
Cs
Earth's crust
6,5∙10 -3
0,03
accurate data
No
The soil
3∙10 -3
0,63
1,36
5∙10 -3
Sea water
1,5∙10 -5
1,06
0,038
2∙10 -5
Plants
1∙10 -5
0,02
5∙10 -4
Animals
10 -4
0,27
10 -5
Human
10 -4
0,08
0,23
10 -5
10 -4
The alkali metals most necessary for the human body are sodium and potassium. Almost all elements enter the human body mainly through food.
Lithium sources.
Lithium is found in some mineral waters, as well as sea and rock salt. It is also found in plants, but its concentration, like any microelements, depends not only on the type and part of the plant, but also on the time of year and even day, on collection conditions and weather, as well as on the area where this plant grows.
In our country, lithium was studied by employees of the Institute of Geochemistry named after Acad. V.I. Vernadsky in Moscow. It was found that the above-ground parts of plants are richer in lithium than the roots. Most lithium is found in plants of the rose family, cloves, and nightshades, which include tomatoes and potatoes. Although within one family the difference in its content can be enormous - several dozen times. This depends on the geographic location and lithium content in the soil.
Sources of sodium.
Sodium is present in various dietary supplements in the form of monosodium glutamate (flavor), sodium saccharin (sweetener), sodium nitrate (preservative), sodium ascorbate (antioxidant) and sodium bicarbonate (baking soda), as well as in some medications (antacids). However, most of the sodium in the diet comes from salt.
NaCl levels are relatively low in all foods that have not been specially processed. However, salt has been used as a preservative and flavoring for several centuries. It is also used as a dye, filler and to control the fermentation process (for example, when baking bread). For this reason, it is added to foods such as ham, sausages, bacon and other meat products, smoked fish and meats, canned vegetables, most butters, margarine, cheese, unsweetened foods, snack foods, and the cereals we eat at home. breakfast.
The recommended sodium intake is 1.5 grams in a day. Excess salt in the diet is associated with an increased likelihood of stomach cancer and is harmful to the kidneys, especially if they have any problems with the urinary system. Excess salt is one of the leading lifestyle factors that leads to hypertension. If hypertension is asymptomatic, it increases the risk of cardiovascular disease and stroke. Current guidelines for the prevention of hypertension have shown that the most effective diet for the prevention and treatment of high blood pressure should be low in sodium and fat and include large amounts of low-fat dairy products (a source of calcium) and fruits and vegetables (a source of potassium). Thus, it is important to change the diet as a whole, rather than focusing on any one component of the diet. Other important positive factors include physical activity and normal body weight.
People with kidney disease and very young children cannot tolerate large amounts of sodium because their kidneys cannot eliminate it. For this reason, you should not add salt to the food of young children.
By law, food labels must list sodium content, but some manufacturers ignore this rule and list the amount of salt.
We remember: “ Table salt can be annoying our health!»
Sources of potassium.
The best source of potassium is plant foods. These are watermelons, melons, oranges, tangerines, bananas, dried fruits (figs, apricots, rose hips). Berries rich in potassium include lingonberries, strawberries, black and red currants. There is a lot of potassium in vegetables (especially potatoes), legumes, wholemeal products, and rice.
The body's reaction to potassium deficiency.
With a lack of potassium in the body, muscle weakness, intestinal lethargy, and cardiac dysfunction are observed.
“I haven’t gotten up yet, I’m already tired” - this is how the doctor figuratively and clearly characterizes potassium deficiency in the body. A low potassium content in the body usually leads to asthenia (mental and physical exhaustion, fatigue), impaired renal function and depletion of the adrenal cortex. There is a risk of disruption of metabolic processes and conductivity in the myocardium.
Potassium deficiency reduces performance, slows down wound healing, and leads to impaired neuromuscular conduction. Dry skin, dullness and weakness of hair are noted (this is a matter of serious concern, especially for women and girls).
Sudden death may occur with increasing stress. There is poor transmission of nerve impulses. Diuretics (diuretics) reduce potassium absorption. When preparing food, it is necessary to pay attention to the fact that potassium compounds are water-soluble. This circumstance requires you to wash products containing it before chopping them and cook them in a small amount of water.
By the way, traditional medicine believes that the passionate desire to drink alcohol is associated with a lack of potassium in the body.
For potassium depletion use potassium chloride KCl 4 - 5 times a day, 1 g.
The body's reaction to excess potassium.
With an excess of potassium in the body, the main functions of the heart are inhibited: a decrease in the excitability of the heart muscle, a slowdown in the heart rate, deterioration of conductivity, and a weakening of the force of heart contractions. In high concentrations, potassium ions cause cardiac arrest in diastole (the contraction phase of the ventricles of the heart). The toxic dose of potassium is 6 g. The lethal dose is 14 g. Potassium salts can be toxic to the body due to the anion associated with the potassium ion, for example, KCN (potassium cyanide).
To regulate the content of these nutrients, you can take into account the data presented in the following table.
4. Practical part
Experience 1.Flame coloring with compounds.
One of the methods for qualitative detection of alkali metal compounds is based on their ability to color the burner flame.
Solutions of alkali metal salts must be poured into test tubes. Wash the iron wire in hydrochloric acid and then ignite it in a burner flame.
Then you need to moisten the wire with a solution of the salt being tested and add it to the flame.
Salts containing lithium cations, as well as lithium color the flames red color, sodium cations and metal sodium- V yellow, potassium cations and metal potassium color the flames violet color. For better observation, you can view the color through blue glass.
Thus, Li +, Na + and K + ions were discovered in solutions of salts LiCl, NaCl, Na 2 CO 3, Na 2 SO 4, NaNO 3, KCl, KNO 3, K 2 CO 3.
Experience 2.Interaction of alkali metals with water.
Add a piece of metal, thoroughly cleaned of the oxide film, into a glass of water. After dissolving the metal, the solution medium was examined using phenolphthalein.
Carry out this experiment with pieces of lithium, sodium and potassium. The reaction with potassium was most active; it was accompanied by the combustion of potassium, violet sparks and gas evolution were observed. Sodium reacted with water, producing yellow sparks, and lithium reacted most calmly.
The resulting solutions with phenolphthalein turned crimson, indicating the presence of alkali in the solution.
2Li + 2H 2 O = 2LiOH + H 2
2Na + 2H 2 O = 2NaOH + H 2
2K + 2H 2 O = 2KOH + H 2
Experience 3. Hydrolysis of sodium and potassium salts.
The nature of the salt solution environment is studied using acid-base indicators.
Universal indicator papers dipped into solutions of alkali metal salts formed by weak acids Na 2 CO 3 and K 2 CO 3 turned blue, which indicates an alkaline reaction of the solutions. hydrolysis occurred in solutions - the interaction of salts with water molecules:
Na 2 CO 3 ↔ 2Na + + CO 3 2-
CO 3 2- + H 2 O ↔ HCO 3 - + OH -
Na 2 CO 3 + H 2 O ↔ NaHCO 3 + NaOH
Solutions of salts of strong acids NaNO 3, KNO 3, NaCl, KCl, LiCl showed a neutral environment (the color of the indicator paper did not change), which means that hydrolysis of these salts does not occur
conclusions
Why is it so important to know the content of chemical elements in the body?
Chemical elements are not synthesized, unlike many organic substances, in the body, but come from outside with food, air, through the skin and mucous membranes. Therefore, the determination of chemical elements allows you to find out about:
how much does your body correspond to the ideal (by the way, about 20% of people do not have any deviations and, thus, live in harmony with nature);
Are you eating right, does your diet provide the necessary set of nutrients;
Do bad habits harm the body?
how safe is the environment in which you live; the food you eat; Your workplace;
do your stomach, intestines, liver, kidneys, skin function well, regulating the processes of absorption and excretion of nutrients;
Do you have any chronic diseases or a predisposition to them?
Are you being treated correctly?
What diseases are most closely related to elemental imbalance?
First of all, this is:
decreased immunity;
diseases of the skin, hair, nails;
scoliosis, osteoporosis, osteochondrosis;
hypertension;
allergies, including bronchial asthma;
diabetes, obesity;
diseases of the cardiovascular system;
blood diseases (anemia);
intestinal dysbiosis, chronic gastritis, colitis;
infertility, decreased potency in men;
impaired growth and development in children.
Many years of experience of doctors shows that more than 80% of the population have a more or less pronounced imbalance of microelements. Therefore, if you have any , you should pay attention to this!
Many scientists believe that not only are all chemical elements present in a living organism, but each of them performs a specific biological function.
We have clarified the biological role of only one group of chemical elements. Alkali metals are extremely important for human health, like most others. It is very important for human health to maintain the optimal concentration of each element: both a deficiency of an element and its excess are harmful.
The stability of the chemical composition of the body is one of the most important and mandatory conditions for its normal functioning. .
There is an erroneous, although widespread, opinion about the possibility of correcting an imbalance in the elemental composition of the human body by enriching the diet with certain products containing the necessary mineral elements. However, it should be taken into account that the presence of necessary macro- and microelements in food and water (which is especially obvious for residents of rural areas) depends to a large extent on the so-called “local biogeochemical cycle” of elements, which determines the content of macro- and microelements in food plants and animals.
A deficiency or excess of certain elements in the human body, as a rule, is a consequence of a deficiency or excess of these elements passing through the food chain: from soil to plants and animals to humans. When a deficiency of any element develops, nutritional correction is not enough, even if products from other regions are used for this purpose, the soils of which are enriched with the necessary microelement.
Only an individual selection of special mineral and other preparations aimed at normalizing the microelement balance of the body will provide real and effective help in the development of a pathological condition.
In conclusion, we present the commandments of traditional and scientific medicine that everyone should know:
Everything is connected to everything.
Everything has to go somewhere.
Nature knows best.
Nothing comes for free.
Used Books
1. Gabrielyan O.S. Chemistry, 9th grade, Textbook for educational institutions. - M. “Bustard”, 2001
2. Glinka N.L. General chemistry, Textbook for universities. - L. “Chemistry”, 1983
3. General chemistry. Chemistry of biogenic elements. Textbook for honey. specialist. call. Yu.A. Ershov and others - M. “Higher School”, 1993
4. Sychev A.P., Fadeev G.N. Chemistry of metals. Tutorial. – M. “Enlightenment”, 1984
5. MHTML. Do c ument. integrated lesson “Alkali metals”. Festival "Open Lesson", 2003
6.
7.
The structure of the outer electronic layers in the atoms of group I elements allows us, first of all, to assume that they do not have a tendency to add electrons. On the other hand, the donation of a single external electron, it would seem, should occur very easily and lead to the formation of stable monovalent cations of the elements in question.
As experience shows, these assumptions are fully justified only in relation to the elements of the left column (Li, Na, K and analogues). For copper and its analogues, they are only half true: in the sense of their lack of tendency to add electrons. At the same time, their 18-electron layer, which is farthest from the nucleus, turns out to be not yet completely fixed and, under certain conditions, is capable of partial loss of electrons. The latter makes it possible to exist, along with monovalent Cu, Agand Aualso compounds of the elements under consideration, corresponding to their higher valency.
Such a discrepancy between assumptions derived from atomic models and experimental results shows that consideration of the properties of elements based ononlythe electronic structures of atoms without taking into account other features are not always sufficient for the chemical characterization of these elements even in the roughest terms.
Alkali metals.
The name alkali metals applied to elements of the Li-Cs series is due to the fact that their hydroxides are strong alkalis. Sodium
And potassium
are among the most common elements, accounting for 2.0 and 1.1%, respectively, of the total number of atoms in the earth’s crust. Contents in it lithium
(0,02%), rubidium
(0.004%) and cesium
(0.00009%) is already significantly less, and France
- negligible. Elementary Na and K were isolated only in 1807. Lithium was discovered in 1817, cesium and rubidium - in 1860 and 1861, respectively. Element No. 87 - francium - was discovered in 1939, and received its name in 1946. Natural sodium and cesium are “pure” elements (23 Na and 133 Cs), lithium is composed of the isotopes 6 Li (7.4%) and 7 Li (92.6%), potassium is made of the isotopes 39 K (93.22%) .
40 K (0.01%) and 41 K (6.77%), rubidium - from the isotopes 85 Rb (72.2%) and 87 Rb (27.8%). Of the isotopes of francium, the most important is the naturally occurring 223 Fr (the average lifespan of an atom is 32 minutes).
Prevalence:
Only compounds of alkali metals are found in nature. Sodium and potassium are permanent constituents of many silicates. Of the individual minerals, sodium is the most important - salt (NaCl) is part of sea water and in certain areas of the earth's surface forms huge deposits of rock salt under a layer of alluvial rocks. The upper layers of such deposits sometimes contain accumulations of potassium salts in the form of layers sylvinite (mKCl∙nNaCl), ka rnallite (KCl MgCl 2 6H 2 O), etc., which serve as the main source for obtaining compounds of this element. Only a few natural accumulations of potassium salts of industrial importance are known. A number of minerals are known for lithium, but their accumulations are rare. Rubidium and cesium occur almost exclusively as impurities in potassium. Traces of France are always contained in uranium ores . Lithium minerals are, for example, spodumene And lepidolite (Li 2 KAl). Part of the potassium in the latter of them is sometimes replaced by rubidium. The same applies to carnallite, which can serve as a good source of rubidium. The relatively rare mineral is most important for cesium technology pollucite - CsAI(SiO 3) 2.
Receipt:
In their free state, alkali metals can be isolated by electrolysis of their molten chloride salts. Sodium is of primary practical importance, the annual world production of which is more than 200 thousand tons. The installation diagram for its production by electrolysis of molten NaCl is shown below. The bath consists of a steel casing with fireclay lining, a graphite anode (A) and an annular iron cathode (K), between which a mesh diaphragm is located. The electrolyte is usually not pure NaCl (mp 800 ℃), but a more fusible mixture of approximately 40% NaCl and 60% CaCl 2, which makes it possible to work at temperatures of about 580 ° C. Metallic sodium, which collects in the upper part of the annular cathode space and passes into the collector, contains a small (up to 5%) admixture of calcium, which is then almost completely released (the solubility of Ca in liquid sodium at its melting point is only 0.01%). As electrolysis progresses, NaCl is added to the bath. Electricity consumption is about 15 kWh per 1 kg Na.
2NaCl→ 2Na+Cl 2
This is interesting:
Before the introduction of the electrolytic method into practice, metallic sodium was obtained by heating soda with coal according to the reaction:
Na 2 CO 3 +2C+244kcal→2Na+3CO
The production of metallic K and Li is incomparably less than that of sodium. Lithium is obtained by electrolysis of the LiCl + KCl melt, and potassium is obtained by the action of sodium vapor on the KCl melt, which flows countercurrently to them in special distillation columns (from the upper part of which potassium vapor comes out). Rubidium and cesium are almost never mined on a large scale. To obtain small quantities of these metals, it is convenient to use heating of their chlorides with metallic calcium in a vacuum.
2LiCl→2Li+Cl 2
Physical properties:
In the absence of air, lithium and its analogs are silvery-white (with the exception of yellowish cesium) substances with a more or less strong metallic luster. All alkali metals are characterized by low densities, low hardness, low melting and boiling points, and good electrical conductivity. Their most important constants are compared below:
|
Density, g/cm3. |
|||||
|
Melting point, °C |
|||||
|
Boiling point, °C |
Due to their low density, Li, Na and K float on water (Li even on kerosene). Alkali metals are easily cut with a knife, and the hardness of the softest of them - cesium - does not exceed the hardness of wax. The non-luminous flame of a gas burner is colored by alkali metals and their volatile compounds in characteristic colors, of which the bright yellow inherent in sodium is the most intense.
This is interesting:
Externally manifested in the form of coloring of the flame, the emission of light rays by heated atoms of alkali metals is caused by the jump of electrons from higher to lower energy levels. For example, the characteristic yellow line in the spectrum of sodium appears when an electron jumps from the 3p level to the 3s level. Obviously, for such a jump to be possible, a preliminary excitation of the atom is necessary, that is, the transfer of one or more of its electrons to a higher energy level. In the case under consideration, excitation is achieved due to the heat of the flame (and requires an expenditure of 48 kcal/g-atom); in general, it can result from the imparting of energy of various types to the atom. Other alkali metals cause the appearance of the following flame colors: Li - carmine-red, K-violet, Rb - bluish-red, Cs - blue.
The luminescence spectrum of the night sky shows the constant presence of yellow sodium radiation. The altitude of the place of its origin is estimated at 200-300 km.T. That is, the atmosphere at these altitudes contains sodium atoms (of course, in negligible quantities). The occurrence of radiation is described by a number of elementary processes (the asterisk indicates the excited state; M is any third particle - O 2, O 0, N 2, etc.): Na + O 0 + M = NaO + M*, then NaO + O=O 2 + Na* and finally Na*= Na +λν.
Sodium and potassium should be stored in tightly closed containers under a layer of dry and neutral kerosene. Their contact with acids, water, chlorinated organic compounds and solid carbon dioxide is unacceptable. Do not accumulate small potassium scraps, which oxidize especially easily (due to their relatively large surface). Unused potassium and sodium residues in small quantities are destroyed by interaction with excess alcohol, in large quantities - by burning on the coals of a fire. Alkali metals that catch fire in a room are best extinguished by covering them with dry soda ash powder.
Chemical properties:
From the chemical point of view, lithium and its analogues are extremely reactive metals (and their activity usually increases in the direction from Li to Cs). In all compounds, alkali metals are monovalent. Located at the extreme left of the voltage series, they energetically interact with water according to the following scheme:
2E + 2H 2 O = 2EON + H 2
When reacting with Li and Na, the release of hydrogen is not accompanied by its ignition; for K it already occurs, and for Rb and Cs the interaction proceeds with an explosion.
· In contact with air, fresh sections of Na and K (to a lesser extent, Li) are immediately covered with a loose film of oxidation products. In view of this, Na and K are usually stored under kerosene. Na and K heated in air easily ignite, while rubidium and cesium spontaneously ignite even at ordinary temperatures.
4E+O 2 →2E 2 O (for lithium)
2E+O 2 →E 2 O 2 (for sodium)
E+O 2 →EO 2(for potassium, rubidium and cesium)
Practical application is mainly found in sodium peroxide (Na 2 0 2). Technically, it is obtained by oxidation at 350°C of atomized sodium metal:
2Na+O 2 →Na 2 O 2 +122kcal
· Melts of simple substances are capable of combining with ammonia to form amides and imides, solvates:
2Na melt +2NH 3 →2NaNH 2 +H 2 (sodium amide)
2Na melt +NH 3 →Na 2 NH+H 2 (sodium imide)
Na melt +6NH 3 → (sodium solvate)
When peroxides interact with water, the following reaction occurs:
2E 2 O 2 +2H 2 O=4EOH+O 2
The interaction of Na 2 O 2 with water is accompanied by hydrolysis:
Na 2 O 2 +2H 2 O→2NaOH + H 2 O 2 +34 kcal
This is interesting:
InteractionNa 2 O 2 with carbon dioxide according to the scheme
2Na 2 O 2 + 2CO 2 =2Na 2 CO 3 +O 2 +111 kcal
serves as the basis for the use of sodium peroxide as a source of oxygen in insulating gas masks and on submarines. Pure or containing various additives (for example, bleach mixed with Ni or C saltsu) sodium peroxide has the technical name "oxylitol". Mixed oxylit preparations are especially convenient for obtaining oxygen, which they release under the influence of water. Oxylitol compressed into cubes can be used to obtain a uniform flow of oxygen in a conventional apparatus for producing gases.
Na 2 O 2 +H 2 O=2NaOH+O 0 (atomic oxygen is released due to the decomposition of hydrogen peroxide).
Potassium superoxide ( KO 2) is often included in oxylitol. Its interaction with carbon dioxide in this case follows the overall equation:
Na 2 O 2 + 2KO 2 + 2CO 2 = Na 2 CO 3 + K 2 CO 3 + 2O 2 + 100 kcal, i.e. carbon dioxide is replaced by an equal volume of oxygen.
· Capable of forming ozonides. The formation of potassium ozonide-KO 3 follows the equation:
4KOH+3O 3 = 4KO 3 + O 2 +2H 2 O
It is a red crystalline substance and is a strong oxidizing agent. During storage, KO 3 decomposes slowly according to the equation 2NaO 3 →2NaO 2 +O 2 +11 kcal already under normal conditions. It instantly decomposes with water according to the overall scheme 4 KO 3 +2 H 2 O=4 KOH +5 O 2
· Capable of reacting with hydrogen to form ionic hydrides, according to the general scheme:
The interaction of hydrogen with heated alkali metals is slower than with alkaline earth metals. In the case of Li, heating to 700-800 °C is required, while its analogs interact already at 350-400 °C. Alkali metal hydrides are very strong reducing agents. Their oxidation by atmospheric oxygen in a dry state is relatively slow, but in the presence of moisture the process accelerates so much that it can lead to spontaneous ignition of the hydride. This especially applies to hydrides K, Rb and Cs. A violent reaction occurs with water according to the following scheme:
EN+ H 2 O= H 2 +EON
EH+O 2 →2EOH
When NaH or KH reacts with carbon dioxide, the corresponding salt of formic acid is formed:
NaH+CO 2 →HCOONa
Capable of forming complexes:
NaH+AlCl 3 →NaAlH 4 +3NaCl (sodium allanate)
NaAlH 4 → NaH+AlH 3
Normal alkali metal oxides (with the exception of Li 2 0) can be prepared only indirectly . They are solids of the following colors:
Na 2 O+2HCl=2NaCl+H 2 O
Alkali metal hydroxides (EOH) are colorless, very hygroscopic substances that corrode most materials that come into contact with them. Hence their sometimes used name in practice - caustic alkalis. When exposed to alkalis, the skin of the human body swells greatly and becomes slippery; with longer action, a very painful deep burn is formed. Caustic alkalis are especially dangerous for the eyes (it is recommended to wear safety glasses when working). Any alkali that gets on your hands or dress should be immediately washed off with water, then the affected area should be moistened with a very dilute solution of any acid and rinsed again with water.
All of them are relatively fusible and volatile without decomposition (except for LiOH, which eliminates water). To obtain hydroxide-alkali metals Electrolytic methods are mainly used. The most large-scale production is sodium hydroxide electrolysis concentrated aqueous table salt solution:
2NaCl+2H 2 O→2NaOH+Cl 2 +H 2
Ø Are typical grounds:
NaOH+HCl=NaCl+H2O
2NaOH+CO 2 =Na 2 CO 3 +H 2 O
2NaOH+2NO 2 =NaNO 3 +NaNO 2 +H 2 O
Ø Capable of forming complexes:
NaOH+ZnCl 2 = (ZnOH)Cl+NaCl
2Al+2NaOH+6H 2 O=2Na+3H 2
Al 2 O 3 + 6NaOH = 2Na 3 AlO 3 + 3H 2 O
Al(OH) 3 +NaOH=Na
Ø Capable of reacting with non-metals:
Cl 2 +2KOH=KCl+KClO+H 2 O (reaction occurs without heating)
Cl 2 +6KOH=5KCl+KClO 3 +3H 2 O (the reaction occurs with heating)
3S+6NaOH=2Na 2 S+Na 2 SO 3 +3H 2 O
Ø Used in organic synthesis (in particular, potassium and sodium hydroxide, sodium hydroxide is indicated in the examples):
NaOH+C 2 H 5 Cl=NaCl+C 2 H 4 (method for producing alkenes, ethylene (ethene) in this case), an alcohol solution of sodium hydroxide was used.
NaOH+C 2 H 5 Cl=NaCl+C 2 H 5 OH(a method for producing alcohols, ethanol in this case), an aqueous solution of sodium hydroxide was used.
2NaOH+C 2 H 5 Cl=2NaCl+C 2 H 2 +H 2 O (method for producing alkynes, acetylene (ethyne) in this case), an alcohol solution of sodium hydroxide was used.
C 6 H 5 OH (phenol) +NaOH= C 6 H 5 ONa+H 2 O
NaOH(+CaO)+CH 3 COONa→Na 2 CO 3 CH 4 (one of the methods for producing methane)
Ø You need to know the decomposition of several salts:
2KNO 3 →2KNO 2 +O 2
4KClO 3→ KCl+3KClO 4
2KClO 3→ KCl+3O 2
4Na 2 SO 3 →Na 2 S+3Na 2 SO 4
It is noteworthy that the decomposition of nitrates occurs approximately in the range of 450-600 ℃, then they melt without decomposition, but when reaching approximately 1000-1500 ℃, decomposition occurs according to the following scheme:
4LiNO 2 →2Li 2 O+4NO+O 2
This is interesting:
K 4 [ Fe(CN) 6 ]+ FeCl 3 = KFe[ Fe(CN) 6 ]+3 KCl(qualitative reaction toFe3+)
3K 4 +4FeCl 3 =Fe 4 3 +12KCl
Na 2 O 2 +2H 2 O=2NaOH+ H 2 O 2
4NaO 2 +2H 2 O=4NaOH+ 3O 2
4NaO 3 +2H 2 O=4NaOH+5O 2 (reaction of sodium ozonide with water )
2NaO 3 → 2NaO 2 +O 2(Decomposition occurs at different temperatures, for example: decomposition of sodium ozonide at -10
°C, cesium ozonide at +100°C)
NaNH 2 +H 2 O→ NaOH+NH 3
Na 2 NH+2H 2 O→ 2NaOH+NH 3
Na 3 N+3H 2 O→3NaOH+NH 3
KNO 2 +2Al+KOH+5H 2 O→2K+NH 3
2NaI + Na 2 O 2 + 2H 2 SO 4 →I 2 ↓+ 2Na 2 SO 4 + 2H 2 O
Fe 3 O 4 +4NaH=4NaOH+3Fe
5NaN 3 +NaNO 3 →8N 2 +3Na 2 O
Application:
Sodium is widely used in the synthesis of organic compounds and, partly, for the preparation of some of its derivatives. In nuclear technology it is used as a coolant.
Lithium is of absolutely exceptional importance for thermonuclear technology. In the rubber industry it is used in the production of artificial rubber (as a polymerization catalyst), in metallurgy - as a valuable additive to some other metals and alloys. For example, adding only hundredths of a percent of lithium greatly increases the hardness of aluminum and its alloys, and adding 0.4% lithium to lead almost triples its hardness without compromising bending resistance. There are indications that a similar cesium additive greatly improves the mechanical properties of magnesium and protects it from corrosion, but this is not the case with its use. Sodium hydride is sometimes used in metallurgy to isolate rare metals from their compounds. Its 2% solution in molten NaOH is used to remove scale from steel products (after a minute of holding in it, the hot product is immersed in water, which is reduced according to the equation
Fe 3 O 4 + 4NaH = 4NaOH + 3Fe (scale disappears).
Schematic diagram of a factory installation for producing soda by ammoniamethod (Solvay, 1863).
Limestone is fired in the furnace (L), and the resulting CO 2 enters the carbonization tower (B), and CaO is quenched with water (C), after which Ca(OH) 2 is pumped into the mixer (D), where it meets NH 4 Cl , this releases ammonia. The latter enters the absorber (D) and saturates a strong NaCl solution there, which is then pumped into the carbonization tower, where, when interacting with CO 2, NaHCO 3 and NH 4 Cl are formed. The first salt is almost completely precipitated and retained on the vacuum filter (E), and the second is pumped back into the mixer (D). Thus, NaCl and limestone are constantly consumed, and NaHCO 3 and CaCl 2 are obtained (the latter in the form of production waste). Sodium bicarbonate is then transferred by heating into soda.
Editor: Galina Nikolaevna Kharlamova
Alkali metals include metals of group IA of the Periodic Table of D.I. Mendeleev - lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr). The outer energy level of alkali metals contains one valence electron. The electronic configuration of the external energy level of alkali metals is ns 1. In their compounds they exhibit a single oxidation state of +1. In OVR they are reducing agents, i.e. give up an electron.
Physical properties of alkali metals
All alkali metals are light (have low density), very soft (with the exception of Li, they are easily cut with a knife and can be rolled into foil), have low boiling and melting points (with an increase in the charge of the nucleus of an alkali metal atom, the melting point decreases).
In the free state, Li, Na, K and Rb are silvery-white metals, Cs is a golden-yellow metal.
Alkali metals are stored in sealed ampoules under a layer of kerosene or petroleum jelly, since they are highly chemically reactive.
Alkali metals have high thermal and electrical conductivity, which is due to the presence of a metallic bond and a body-centered crystal lattice
Preparation of alkali metals
All alkali metals can be obtained by electrolysis of the melt of their salts, but in practice only Li and Na are obtained in this way, which is associated with the high chemical activity of K, Rb, Cs:
2LiCl = 2Li + Cl 2
2NaCl = 2Na + Cl2
Any alkali metal can be obtained by reducing the corresponding halide (chloride or bromide), using Ca, Mg or Si as reducing agents. Reactions are carried out with heating (600 – 900C) and under vacuum. The general equation for obtaining alkali metals in this way is:
2MeCl + Ca = 2Me + CaCl 2,
where Me is a metal.
There is a known method for producing lithium from its oxide. The reaction is carried out by heating to 300°C and under vacuum:
2Li 2 O + Si + 2CaO = 4Li + Ca 2 SiO 4
Potassium can be produced by the reaction between molten potassium hydroxide and liquid sodium. The reaction is carried out by heating to 440°C:
KOH + Na = K + NaOH
Chemical properties of alkali metals
All alkali metals actively interact with water forming hydroxides. Due to the high chemical activity of alkali metals, the reaction of interaction with water may be accompanied by an explosion. Lithium reacts most calmly with water. The general reaction equation is:
2Me + H2O = 2MeOH + H2
where Me is a metal.
Alkali metals interact with atmospheric oxygen to form a number of different compounds - oxides (Li), peroxides (Na), superoxides (K, Rb, Cs):
4Li + O 2 = 2Li 2 O
2Na + O 2 = Na 2 O 2
All alkali metals react with nonmetals (halogens, nitrogen, sulfur, phosphorus, hydrogen, etc.) when heated. For example:
2Na + Cl 2 = 2NaCl
6Li + N 2 = 2Li 3 N
2Li +2C = Li 2 C 2
2Na + H 2 = 2NaH
Alkali metals are capable of interacting with complex substances (acid solutions, ammonia, salts). Thus, when alkali metals interact with ammonia, amides are formed:
2Li + 2NH 3 = 2LiNH 2 + H 2
The interaction of alkali metals with salts occurs according to the following principle - they displace less active metals (see the activity series of metals) from their salts:
3Na + AlCl 3 = 3NaCl + Al
The interaction of alkali metals with acids is ambiguous, since when such reactions occur, the metal will initially react with the water of the acid solution, and the alkali formed as a result of this interaction will react with the acid.
Alkali metals react with organic substances, such as alcohols, phenols, carboxylic acids:
2Na + 2C 2 H 5 OH = 2C 2 H 5 ONa + H 2
2K + 2C 6 H 5 OH = 2C 6 H 5 OK + H 2
2Na + 2CH 3 COOH = 2CH 3 COONa + H 2
Qualitative reactions
A qualitative reaction to alkali metals is the coloring of the flame by their cations: Li + colors the flame red, Na + yellow, and K + , Rb + , Cs + purple.
Examples of problem solving
EXAMPLE 1
| Exercise | Carry out the chemical transformations Na→Na 2 O→NaOH→Na 2 SO 4 | |||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| Solution | 4Na + O 2 →2Na 2 O The most active among metals are alkali metals. They actively react with simple and complex substances. General informationAlkali metals are in group I of the periodic table. These are soft monovalent metals of gray-silver color with a low melting point and low density. They exhibit a single oxidation state of +1, being reducing agents. Electronic configuration - ns 1.
Rice. 1. Sodium and lithium. The general characteristics of group I metals are given in the table.
Active metals react quickly with other substances, so they are found in nature only in minerals. ReceiptSeveral methods are used to obtain pure alkali metal:
electrolysis of melts, most often chlorides or hydroxides - 2NaCl → 2Na + Cl 2, 4NaOH → 4Na + 2H 2 O + O 2; calcination of soda (sodium carbonate) with coal to obtain sodium - Na 2 CO 3 + 2C → 2Na + 3CO; reduction of rubidium from chloride by calcium at high temperatures - 2RbCl + Ca → 2Rb + CaCl 2 ; InteractionThe properties of alkali metals are determined by their structure. Being in the first group of the periodic table, they have only one valence electron in the outer energy level. A single electron easily goes to the oxidizing atom, which contributes to the rapid entry into the reaction. Metallic properties increase in the table from top to bottom, so lithium loses its valence electron more easily than francium. Lithium is the hardest element among all alkali metals. The reaction of lithium with oxygen occurs only under the influence of high temperature. Lithium reacts with water much more slowly than the other metals in the group. General chemical properties are presented in the table.
With a high-quality reaction, they have different flame colors. Lithium burns with a crimson flame, sodium with a yellow flame, and cesium with a pink-violet flame. Alkali metal oxides also have different colors. Sodium turns white, rubidium and potassium turn yellow.
Rice. 2. Qualitative reaction of alkali metals. ApplicationSimple metals and their compounds are used to make light alloys, metal parts, fertilizers, soda and other substances. Rubidium and potassium are used as catalysts. Sodium vapor is used in fluorescent lamps. Only francium has no practical use due to its radioactive properties. How group I elements are used is briefly described in the table on the use of alkali metals.
Rice. 3. Baking soda. What have we learned?From the 9th grade lesson we learned about the features of alkali metals. They are in group I of the periodic table and give up one valence electron during reactions. These are soft metals that easily enter into chemical reactions with simple and complex substances - halogens, non-metals, acids, water. They are found in nature only as part of other substances, so electrolysis or a reduction reaction is used to extract them. They are used in industry, construction, metallurgy, and energy. Test on the topicEvaluation of the reportAverage rating: 4.4. Total ratings received: 91. |
