Atomic mass of sodium. Sodium. Properties of sodium. Use of sodium. Examples of problem solving

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SODIUM– (Natrium) Na, a chemical element of group 1 (Ia) of the Periodic Table, belongs to the alkaline elements. Atomic number 11, relative atomic mass 22.98977. In nature there is one stable isotope 23 Na. Six radioactive isotopes of this element are known, two of which are of interest to science and medicine. Sodium-22, with a half-life of 2.58 years, is used as a source of positrons. Sodium-24 (its half-life is about 15 hours) is used in medicine for the diagnosis and treatment of some forms of leukemia.

Oxidation state +1.

Sodium compounds have been known since ancient times. Sodium chloride is an essential component of human food. It is believed that people began to use it in the Neolithic, i.e. about 5–7 thousand years ago.

The Old Testament mentions a substance called “neter.” This substance was used as a detergent. Most likely, neter is soda, a sodium carbonate that formed in the salty Egyptian lakes with calcareous shores. The Greek authors Aristotle and Dioscorides later wrote about the same substance, but under the name “nitron,” and the ancient Roman historian Pliny the Elder, mentioning the same substance, called it “nitrum.”

In the 18th century Chemists already knew a lot of different sodium compounds. Sodium salts were widely used in medicine, in tanning leather, and in dyeing fabrics.

Metallic sodium was first obtained by the English chemist and physicist Humphry Davy by electrolysis of molten sodium hydroxide (using a voltaic column of 250 pairs of copper and zinc plates). The name "sodium" chosen by Davy for this element reflects its origin from the soda Na 2 CO 3 . Latin and Russian names element are derived from the Arabic “natrun” (natural soda).

Distribution of sodium in nature and its industrial extraction.

Sodium is the seventh most abundant element and the fifth most abundant metal (after aluminum, iron, calcium and magnesium). Its contents in earth's crust is 2.27%. Most of the sodium is found in various aluminosilicates.

Huge deposits of sodium salts in relatively pure form exist on all continents. They are the result of the evaporation of ancient seas. This process is still ongoing in Salt Lake (Utah), the Dead Sea and other places. Sodium is found in the form of NaCl chloride (halite, rock salt), as well as carbonate Na 2 CO 3 NaHCO 3 2H 2 O (trona), nitrate NaNO 3 (saltpeter), sulfate Na 2 SO 4 10H 2 O (mirabilite) , tetraborate Na 2 B 4 O 7 10 H 2 O (borax) and Na 2 B 4 O 7 4H 2 O (kernite) and other salts.

There are inexhaustible reserves of sodium chloride in natural brines and ocean waters (about 30 kg m–3). It is estimated that rock salt in an amount equivalent to the sodium chloride content in the World Ocean would occupy a volume of 19 million cubic meters. km (50% more than the total volume of the North American continent above sea level). A prism of this volume with a base area of 1 sq. km can reach the Moon 47 times.

Currently, the total production of sodium chloride from sea water reached 6–7 million tons per year, which is about a third of total world production.

Living matter contains an average of 0.02% sodium; There is more of it in animals than in plants.

Characteristics of a simple substance and industrial production of sodium metal.

Sodium is a silvery-white metal, in thin layers with a violet tint, plastic, even soft (easily cut with a knife), a fresh cut of sodium is shiny. The electrical and thermal conductivities of sodium are quite high, the density is 0.96842 g/cm 3 (at 19.7° C), the melting point is 97.86° C, and the boiling point is 883.15° C.

The ternary alloy, containing 12% sodium, 47% potassium and 41% cesium, has the lowest melting point for metal systems, equal to –78 ° C.

Sodium and its compounds color the flame bright yellow. The double line in the sodium spectrum corresponds to transition 3 s 1–3p 1 in the atoms of the element.

The chemical activity of sodium is high. In air, it quickly becomes covered with a film of a mixture of peroxide, hydroxide and carbonate. Sodium burns in oxygen, fluorine and chlorine. When a metal is burned in air, Na 2 O 2 peroxide is formed (with an admixture of Na 2 O oxide).

Sodium reacts with sulfur when ground in a mortar and reduces sulfuric acid to sulfur or even sulfide. Solid carbon dioxide (“dry ice”) explodes on contact with sodium ( carbon dioxide fire extinguishers Sodium cannot be used to extinguish a fire!). With nitrogen, the reaction occurs only in an electrical discharge. Sodium does not interact only with inert gases.

Sodium reacts actively with water:

2Na + 2H 2 O = 2NaOH + H 2

The heat released during the reaction is enough to melt the metal. Therefore, if small piece throw sodium into water, it melts due to the thermal effect of the reaction and a drop of metal, which is lighter than water, “runs” along the surface of the water, driven by the reactive force of the released hydrogen. Sodium reacts much more calmly with alcohols than with water:

2Na + 2C 2 H 5 OH = 2C 2 H 5 ONa + H 2

Sodium readily dissolves in liquid ammonia to form bright blue metastable solutions with unusual properties. At –33.8° C, up to 246 g of sodium metal dissolves in 1000 g of ammonia. Dilute solutions are blue, concentrated solutions are bronze. They can be stored for about a week. It has been established that in liquid ammonia, sodium ionizes:

Na Na + + e –

The equilibrium constant of this reaction is 9.9·10 –3. The leaving electron is solvated by ammonia molecules and forms a complex –. The resulting solutions have metallic electrical conductivity. When ammonia evaporates, the original metal remains. At long-term storage solution, it gradually becomes discolored due to the reaction of the metal with ammonia to form the amide NaNH 2 or imide Na 2 NH and the release of hydrogen.

Store sodium under a layer of dehydrated liquid (kerosene, mineral oil), transported only in sealed metal containers.

The electrolytic method for the industrial production of sodium was developed in 1890. The melt was subjected to electrolysis caustic soda, as in Davy’s experiments, but using more advanced energy sources than a voltaic column. In this process, along with sodium, oxygen is released:

anode (nickel): 4OH – – 4e – = O 2 + 2H 2 O.

During the electrolysis of pure sodium chloride, serious problems arise, associated, firstly, with the close melting point of sodium chloride and the boiling point of sodium and, secondly, with the high solubility of sodium in liquid sodium chloride. Adding potassium chloride, sodium fluoride, calcium chloride to sodium chloride allows you to reduce the melt temperature to 600° C. Production of sodium by electrolysis of a molten eutectic mixture (an alloy of two substances with the lowest melting point) 40% NaCl and 60% CaCl 2 at ~580° C in a cell developed by the American engineer G. Downs, it was started in 1921 by DuPont near the power plant at Niagara Falls.

The following processes occur on the electrodes:

cathode (iron): Na + + e – = Na

Ca 2+ + 2e – = Ca

anode (graphite): 2Cl – – 2e – = Cl 2.

Sodium and calcium metals form on a cylindrical steel cathode and are lifted up by a cooled tube in which the calcium solidifies and falls back into the melt. Chlorine generated at the central graphite anode is collected under the nickel roof and then purified.

Currently, the production volume of sodium metal is several thousand tons per year.

The industrial use of sodium metal is due to its strong reducing properties. For a long time, most of the metal produced was used to produce tetraethyl lead PbEt 4 and tetramethyl lead PbMe 4 (anti-knock agents for gasoline) by the reaction of alkyl chlorides with an alloy of sodium and lead at high blood pressure. Now this production is rapidly declining due to environmental pollution.

Another area of application is the production of titanium, zirconium and other metals by reducing their chlorides. Smaller amounts of sodium are used to produce compounds such as hydride, peroxide and alcoholates.

Dispersed sodium is a valuable catalyst in the production of rubber and elastomers.

There is increasing use of molten sodium as a heat exchange fluid in fast neutron nuclear reactors. Sodium's low melting point, low viscosity, small neutron absorption cross section, combined with extremely high heat capacity and thermal conductivity, make it (and its alloys with potassium) an indispensable material for these purposes.

Sodium reliably cleans transformer oils, ethers and other organic substances from traces of water, and with the help of sodium amalgam you can quickly determine the moisture content in many compounds.

Sodium compounds.

Sodium forms a complete set of compounds with all the usual anions. It is believed that in such compounds there is almost complete separation of charge between the cationic and anionic parts of the crystal lattice.

Sodium oxide Na 2 O is synthesized by the reaction of Na 2 O 2, NaOH, and most preferably NaNO 2, with sodium metal:

Na 2 O 2 + 2Na = 2Na 2 O

2NaOH + 2Na = 2Na2O + H2

2NaNO 2 + 6Na = 4Na 2 O + N 2

In the last reaction, sodium can be replaced with sodium azide NaN 3:

5NaN3 + NaNO2 = 3Na2O + 8N2

It is best to store sodium oxide in anhydrous gasoline. It serves as a reagent for various syntheses.

Sodium peroxide Na 2 O 2 in the form of a pale yellow powder is formed by the oxidation of sodium. In this case, under conditions of limited supply of dry oxygen (air), Na 2 O oxide is first formed, which then turns into Na 2 O 2 peroxide. In the absence of oxygen, sodium peroxide is thermally stable up to ~675°C.

Sodium peroxide is widely used in industry as a bleaching agent for fibers, paper pulp, wool, etc. It is a strong oxidizing agent: it explodes when mixed with aluminum powder or charcoal, reacts with sulfur (and becomes hot), and ignites many organic liquids. Sodium peroxide reacts with carbon monoxide to form carbonate. The reaction of sodium peroxide with carbon dioxide releases oxygen:

2Na 2 O 2 + 2CO 2 = 2Na 2 CO 3 + O 2

This reaction is important practical application in breathing apparatus for submariners and firefighters.

Sodium superoxide NaO 2 is obtained by slowly heating sodium peroxide at 200–450° C under an oxygen pressure of 10–15 MPa. Evidence of the formation of NaO 2 was first obtained in the reaction of oxygen with sodium dissolved in liquid ammonia.

The action of water on sodium superoxide leads to the release of oxygen even in the cold:

2NaO 2 + H 2 O = NaOH + NaHO 2 + O 2

As the temperature rises, the amount of oxygen released increases, as the resulting sodium hydroperoxide decomposes:

4NaO 2 + 2H 2 O = 4NaOH + 3O 2

Sodium superoxide is a component of systems for air regeneration in confined spaces.

Sodium ozonide NaO 3 is formed by the action of ozone on anhydrous sodium hydroxide powder at low temperature, followed by extraction of red NaO 3 with liquid ammonia.

Sodium hydroxide NaOH is often called caustic soda or caustic soda. This is a strong base and is classified as a typical alkali. Numerous NaOH hydrates have been obtained from aqueous solutions of sodium hydroxide n H 2 O, where n= 1, 2, 2.5, 3.5, 4, 5.25 and 7.

Sodium hydroxide is very aggressive. It destroys glass and porcelain by interacting with the silicon dioxide they contain:

2NaOH + SiO 2 = Na 2 SiO 3 + H 2 O

The name "caustic soda" reflects the corrosive effect of sodium hydroxide on living tissue. Getting this substance into the eyes is especially dangerous.

The Duke of Orleans' physician, Nicolas Leblanc (1742–1806), developed a convenient process for producing sodium hydroxide from NaCl in 1787 (patent 1791). This first large-scale industrial chemical process was a major technological achievement in Europe in the 19th century. The Leblanc process was later replaced by the electrolytic process. In 1874, world production of sodium hydroxide amounted to 525 thousand tons, of which 495 thousand tons were obtained by the Leblanc method; by 1902, the production of sodium hydroxide reached 1800 thousand tons, but only 150 thousand tons were obtained using the Leblanc method.

Today, sodium hydroxide is the most important alkali in industry. Annual production in the USA alone exceeds 10 million tons. It is obtained in huge quantities electrolysis of brines. When a solution of sodium chloride is electrolyzed, sodium hydroxide is formed and chlorine is released:

cathode (iron) 2H 2 O + 2 e– = H 2 + 2OH –

anode (graphite) 2Cl – – 2 e– = Cl 2

Electrolysis is accompanied by the concentration of alkali in huge evaporators. The largest in the world (at the PPG Inductries" Lake Charles plant) has a height of 41 m and a diameter of 12 m. About half of the sodium hydroxide produced is used directly in the chemical industry to produce various organic and inorganic substances: phenol, resorcinol, b-naphthol, sodium salts (hypochlorite, phosphate, sulfide, aluminates). In addition, sodium hydroxide is used in the production of paper and pulp, soap and detergents, oils, textiles. It is also necessary when processing bauxite. An important application of sodium hydroxide is the neutralization of acids.

Sodium chloride NaCl is known as table salt, rock salt. It forms colorless, slightly hygroscopic cubic crystals. Sodium chloride melts at 801° C, boils at 1413° C. Its solubility in water depends little on temperature: 35.87 g of NaCl dissolves in 100 g of water at 20° C, and 38.12 g at 80° C.

Sodium chloride is a necessary and indispensable food seasoning. In the distant past, salt was equal in price to gold. In ancient Rome, legionnaires were often paid not in money, but in salt, hence the word soldier.

In Kievan Rus they used salt from the Carpathian region, from salt lakes and estuaries on the Black and Seas of Azov. It was so expensive that at ceremonial feasts it was served on the tables of noble guests, while others went away “slurping unsalted.”

After the annexation of the Astrakhan region to the Moscow state, the Caspian lakes became important sources of salt, and still there was not enough of it, it was expensive, so there was discontent among the poorest segments of the population, which grew into an uprising known as the Salt Riot (1648)

In 1711 Peter I issued a decree introducing a salt monopoly. Trade in salt became the exclusive right of the state. The salt monopoly lasted for more than a hundred and fifty years and was abolished in 1862.

Nowadays sodium chloride is a cheap product. Together with coal, limestone and sulfur, it is one of the so-called “big four” mineral raw materials, the most essential for the chemical industry.

Most sodium chloride is produced in Europe (39%), North America (34%) and Asia (20%), while South America and Oceania account for only 3% each, and Africa – 1%. Rock salt forms vast underground deposits (often hundreds of meters thick) that contain more than 90% NaCl. A typical Cheshire salt deposit (the main source of sodium chloride in Great Britain) covers an area of 60 × 24 km and has a salt bed of about 400 m thick. This deposit alone is estimated to be worth more than 10 11 tons.

World salt production by the beginning of the 21st century. reached 200 million tons, 60% of which is consumed by the chemical industry (for the production of chlorine and sodium hydroxide, as well as paper pulp, textiles, metals, rubbers and oils), 30% by the food industry, 10% by other fields of activity. Sodium chloride is used, for example, as a cheap deicing agent.

Sodium carbonate Na 2 CO 3 is often called soda ash or simply soda. It is found in nature in the form of ground brines, brine in lakes and the minerals natron Na 2 CO 3 ·10H 2 O, thermonatrite Na 2 CO 3 ·H 2 O, trona Na 2 CO 3 ·NaHCO 3 ·2H 2 O. Sodium forms and other various hydrated carbonates, bicarbonates, mixed and double carbonates, for example Na 2 CO 3 7H 2 O, Na 2 CO 3 3NaHCO 3, aKCO 3 n H 2 O, K 2 CO 3 NaHCO 3 2H 2 O.

Among the salts of alkali elements obtained industrially, sodium carbonate is of greatest importance. Most often, the method developed by the Belgian chemist-technologist Ernst Solvay in 1863 is used for its production.

A concentrated aqueous solution of sodium chloride and ammonia is saturated with carbon dioxide under slight pressure. In this case, a precipitate of relatively poorly soluble sodium bicarbonate is formed (the solubility of NaHCO 3 is 9.6 g per 100 g of water at 20 ° C):

NaCl + NH 3 + H 2 O + CO 2 = NaHCO 3 Ї + NH 4 Cl

To obtain soda, sodium bicarbonate is calcined:

The carbon dioxide released is returned to the first process. Additional carbon dioxide is obtained by calcining calcium carbonate (limestone):

The second product of this reaction, calcium oxide (lime), is used to regenerate ammonia from ammonium chloride:

Thus, the only by-product of soda production using the Solvay method is calcium chloride.

Overall process equation:

2NaCl + CaCO 3 = Na 2 CO 3 + CaCl 2

Obviously, under normal conditions in an aqueous solution the reverse reaction occurs, since the equilibrium in this system is completely shifted from right to left due to the insolubility of calcium carbonate.

Soda ash obtained from natural raw materials (natural soda ash) is of better quality compared to soda produced by the ammonia method (chloride content less than 0.2%). In addition, specific capital investments and the cost of soda from natural raw materials are 40–45% lower than those obtained synthetically. About a third of the world's soda production now comes from natural deposits.

World production of Na 2 CO 3 in 1999 was distributed as follows:

Total
North America
Asia/Oceania
Zap. Europe
East Europe
Africa
Lat. America

The world's largest producer of natural soda ash is the USA, where the largest proven reserves of trona and brine of soda lakes are concentrated. The deposit in Wyoming forms a layer 3 m thick and an area of 2300 km 2. Its reserves exceed 10 10 tons. In the USA, the soda industry is focused on natural raw materials; the last soda synthesis plant was closed in 1985. Production of soda ash in the United States has stabilized at 10.3–10.7 million tons in recent years.

Unlike the United States, most countries in the world depend almost entirely on the production of synthetic soda ash. China ranks second in the world in soda ash production after the United States. The production of this chemical in China in 1999 reached approximately 7.2 million tons. The production of soda ash in Russia in the same year amounted to about 1.9 million tons.

In many cases, sodium carbonate is interchangeable with sodium hydroxide (for example, in the production of paper pulp, soap, cleaning products). About half of the sodium carbonate is used in the glass industry. One growing application is the removal of sulfur contaminants from gas emissions from power generation plants and large furnaces. Sodium carbonate powder is added to the fuel, which reacts with sulfur dioxide to form solid products, particularly sodium sulfite, which can be filtered or precipitated.

Sodium carbonate was previously widely used as "washing soda", but this application has now disappeared due to the use of other household detergents.

Sodium bicarbonate NaHCO 3 ( baking soda), is used mainly as a source of carbon dioxide in the baking of bread, the manufacture of confectionery, the production of carbonated drinks and artificial mineral waters, as a component of fire extinguishing compositions and medicine. This is due to the ease of its decomposition at 50–100° C.

Sodium sulfate Na 2 SO 4 occurs in nature in anhydrous form (thenardite) and in the form of decahydrate (mirabilite, Glauber's salt). It is part of astrachonite Na 2 Mg(SO 4) 2 4H 2 O, vanthoffite Na 2 Mg(SO 4) 2, glauberite Na 2 Ca(SO 4) 2. The largest reserves of sodium sulfate are in the CIS countries, as well as in the USA, Chile, and Spain. Mirabilite, isolated from natural deposits or brine of salt lakes, is dehydrated at 100 ° C. Sodium sulfate is also a by-product of the production of hydrogen chloride using sulfuric acid, as well as the final product hundreds of industrial processes that use neutralization of sulfuric acid with sodium hydroxide.

Data on the production of sodium sulfate are not published, but global production of the natural raw material is estimated to be about 4 million tons per year. The recovery of sodium sulfate as a by-product is estimated globally at 1.5–2.0 million tons.

For a long time, sodium sulfate was little used. Now this substance is the basis of the paper industry, since Na 2 SO 4 is the main reagent in kraft pulping for the preparation of brown wrapping paper and corrugated cardboard. Wood shavings or sawdust are processed in a hot alkaline solution of sodium sulfate. It dissolves lignin (the component of wood that holds the fibers together) and releases the cellulose fibers, which are then sent to paper making machines. The remaining solution is evaporated until it is capable of burning, providing steam for the plant and heat for evaporation. Molten sodium sulfate and hydroxide are flame resistant and can be reused.

A smaller portion of sodium sulfate is used in the production of glass and detergents. The hydrated form of Na 2 SO 4 ·10H 2 O (Glauber's salt) is a laxative. It is used less now than before.

Sodium nitrate NaNO 3 is called sodium or Chilean nitrate. The large deposits of sodium nitrate found in Chile appear to have been formed by the biochemical decomposition of organic remains. The ammonia released initially was probably oxidized to nitrogen and nitric acids, which then reacted with dissolved sodium chloride.

Sodium nitrate is obtained by the absorption of nitrous gases (a mixture of nitrogen oxides) with a solution of sodium carbonate or hydroxide, or by the exchange interaction of calcium nitrate with sodium sulfate.

Sodium nitrate is used as a fertilizer. It is a component of liquid salt refrigerants, quenching baths in the metalworking industry, and heat-storing compositions. A ternary mixture of 40% NaNO 2, 7% NaNO 3 and 53% KNO 3 can be used from the melting point (142° C) to ~600° C. Sodium nitrate is used as an oxidizing agent in explosives, rocket fuels, and pyrotechnic compositions. It is used in the production of glass and sodium salts, including nitrite, which serves as a food preservative.

Sodium nitrite NaNO 2 can be obtained by thermal decomposition of sodium nitrate or its reduction:

NaNO 3 + Pb = NaNO 2 + PbO

For industrial production sodium nitrite absorbs nitrogen oxides with an aqueous solution of sodium carbonate.

Sodium nitrite NaNO 2, in addition to being used with nitrates as heat-conducting melts, is widely used in the production of azo dyes, for corrosion inhibition and meat preservation.

Elena Savinkina

Atomic mass is the sum of the masses of all protons, neutrons and electrons that make up an atom or molecule. Compared to protons and neutrons, the mass of electrons is very small, so it is not taken into account in calculations. Although this is incorrect from a formal point of view, it is often this term used to indicate the average atomic mass of all isotopes of an element. This is actually relative atomic mass, also called atomic weight element. Atomic weight is the average of the atomic masses of all isotopes of an element found in nature. Chemists must differentiate between these two types of atomic mass when doing their work—an incorrect atomic mass value can, for example, result in an incorrect result for the yield of a reaction product.

Steps

Finding atomic mass from the periodic table of elements

Learn how atomic mass is written. Atomic mass, that is, the mass of a given atom or molecule, can be expressed in standard SI units - grams, kilograms, and so on. However, because atomic masses expressed in these units are extremely small, they are often written in unified atomic mass units, or amu for short. – atomic mass units. One atomic mass unit is equal to 1/12 the mass of the standard isotope carbon-12.

The atomic mass unit characterizes the mass one mole of a given element in grams. This quantity is very useful in practical calculations, since it can be used to easily convert the mass of a given number of atoms or molecules of a given substance into moles, and vice versa.

Find the atomic mass in the periodic table. Most standard periodic tables contain the atomic masses (atomic weights) of each element. Typically, they are listed as a number at the bottom of the element cell, below the letters representing the chemical element. Usually this is not a whole number, but a decimal fraction.

Remember that the periodic table gives the average atomic masses of elements. As noted earlier, the relative atomic masses given for each element in the periodic table are the average of the masses of all isotopes of the atom. This average value is valuable for many practical purposes: for example, it is used in calculating the molar mass of molecules consisting of several atoms. However, when you are dealing with individual atoms, this value is usually not enough.
- Since the average atomic mass is an average of several isotopes, the value shown in the periodic table is not accurate the value of the atomic mass of any single atom.
- The atomic masses of individual atoms must be calculated taking into account exact number protons and neutrons in a single atom.
Calculation of the atomic mass of an individual atom
1. Find the atomic number of a given element or its isotope. Atomic number is the number of protons in the atoms of an element and never changes. For example, all hydrogen atoms, and only they have one proton. The atomic number of sodium is 11 because it has eleven protons in its nucleus, while the atomic number of oxygen is eight because it has eight protons in its nucleus. You can find the atomic number of any element in the periodic table - in almost all its standard versions, this number is indicated above the letter designation of the chemical element. The atomic number is always a positive integer.
  - Suppose we are interested in the carbon atom. Carbon atoms always have six protons, so we know that its atomic number is 6. In addition, we see that in the periodic table, at the top of the cell with carbon (C) is the number "6", indicating that the atomic carbon number is six.
  - Note that the atomic number of an element is not uniquely related to its relative atomic mass in the periodic table. Although, especially for the elements at the top of the table, it may appear that an element's atomic mass is twice its atomic number, it is never calculated by multiplying the atomic number by two.
2. Find the number of neutrons in the nucleus. The number of neutrons can be different for different atoms of the same element. When two atoms of the same element with the same number of protons have different numbers of neutrons, they are different isotopes of that element. Unlike the number of protons, which never changes, the number of neutrons in the atoms of a given element can often change, so the average atomic mass of an element is written as a decimal fraction with a value lying between two adjacent whole numbers.
  
  Add up the number of protons and neutrons. This will be the atomic mass of this atom. Ignore the number of electrons that surround the nucleus - their total mass is extremely small, so they have virtually no effect on your calculations.
Calculating the relative atomic mass (atomic weight) of an element
1. Determine which isotopes are contained in the sample. Chemists often determine the isotope ratios of a particular sample using a special instrument called a mass spectrometer. However, in training, this data will be provided to you in assignments, tests, and so on in the form of values taken from the scientific literature.
  - In our case, let's say that we are dealing with two isotopes: carbon-12 and carbon-13.
2. Determine the relative abundance of each isotope in the sample. For each element, different isotopes occur in different ratios. These ratios are almost always expressed as percentages. Some isotopes are very common, while others are very rare—sometimes so rare that they are difficult to detect. These values can be determined using mass spectrometry or found in a reference book.
  - Let's assume that the concentration of carbon-12 is 99% and carbon-13 is 1%. Other carbon isotopes really exist, but in quantities so small that in this case they can be neglected.
3. Multiply the atomic mass of each isotope by its concentration in the sample. Multiply the atomic mass of each isotope by its percentage abundance (expressed as a decimal). To convert interest to decimal, simply divide them by 100. The resulting concentrations should always add up to 1.
  - Our sample contains carbon-12 and carbon-13. If carbon-12 makes up 99% of the sample and carbon-13 makes up 1%, then multiply 12 (the atomic mass of carbon-12) by 0.99 and 13 (the atomic mass of carbon-13) by 0.01.
  - The reference books give percentages based on the known quantities of all isotopes of a particular element. Most chemistry textbooks contain this information in a table at the end of the book. For the sample being studied, the relative concentrations of isotopes can also be determined using a mass spectrometer.
4. Add up the results. Sum up the multiplication results you got in the previous step. As a result of this operation, you will find the relative atomic mass of your element - the average value of the atomic masses of the isotopes of the element in question. When considering an element as a whole, rather than a specific isotope of a given element, this is the value used.
  - In our example, 12 x 0.99 = 11.88 for carbon-12, and 13 x 0.01 = 0.13 for carbon-13. The relative atomic mass in our case is 11.88 + 0.13 = 12,01 .
- Some isotopes are less stable than others: they break down into atoms of elements with fewer protons and neutrons in the nucleus, releasing particles that make up the atomic nucleus. Such isotopes are called radioactive.

Early 19th century sodium called sodium. This name was given to the element by Humphry Davy, who managed to isolate the metal from alkali. The chemist slightly moistened it and subjected it to electrolysis. Wilhelm Hilbert suggested changing the name of the element discovered by Humphrey.

This is the author of the famous “Annals of Physics”. The scientist began to call sodium sodium. The work was continued by Jens Berzelius. A chemist from Sweden introduced the abbreviation sodium. Our next material is about the properties and role of this metal in people’s lives.

Chemical and physical properties sodium

The element is included in the main subgroup of the 1st group, occupying the 11th position. All metals in the column are alkaline, so sodium. Water reacts with him. A caustic alkali is formed. Another common feature of the metals of the group is the presence of only 1 electron in the outer orbit of the atom.

This makes sodium a powerful reducing agent. An element easily donates an outer-level electron, increasing its oxidation state. The atom comes to the completed shell of the previous level.

Reducing activity is the reason for the absence of pure metal in nature. You can only find connections. So, sodium chloride- This is table salt. Sodium carbonate- baking soda. So, it was not for nothing that the metal was originally called sodium.

It could also be called neter. It is under this name that the substance appears in the Old Testament. As for, people began to use it back in Paleolithic times, that is, about 6,000 years ago.

Most energetic sodium interacts not only with water, but also with all non-metals. To obtain a delicate one from an active reducing agent, the 11th element is combined with. An amalgam is formed.

If, on the contrary, it is necessary to increase the aggression of sodium, oxygen is added to it. The result is peroxide, a powerful oxidizing agent. In its atmosphere, most substances ignite.

Difficulty and “character” sodium hydroxide. It is called caustic. The compound corrodes fabrics, skin, and other organics and materials made on its basis. True, in the air sodium hydroxide captures carbon dioxide and is neutralized, turning into carbonate.

When Humphry Davy managed to isolate pure sodium, the world learned what he was like externally. The metal is silvery white. Thin sections have a purple tint. Sodium formula makes it soft and pliable.

It can be cut with a regular knife and shines on the sliding surface of the blade. The element has a low melting point - 97 degrees. Sodium boils at 883 on the Celsius scale.

Pure metal conducts current and heat well and is not very dense. Its indicator for the 11th element is less than one. This does not discount the weight of sodium in biological processes.

Metal is found in both plants and animal tissues. So, in the human body sodium solution– part of blood plasma, digestive and lymph.

Osmotic pressure, for example, is maintained precisely by the 11th element. It was used not only by nature, but also by man. Mass of sodium goes, for example, for industrial purposes. Which? We'll talk about this in the next chapter.

Applications of sodium

In nature, the element is represented by only one stable isotope. Its mass number is 23. V artificial conditions 15 more isotopes were created. They are not stable, most are destroyed in a couple of minutes. Exception: - 22nd and 24th Na.

The half-lives of these isotopes are significant. The first one lasts for 2.5 years, actively emits positrons, which serves science. The half-life of the second is 15 hours. Enough to serve medicine and help leukemia patients.

In the field of nuclear energy, sodium has become a coolant. The “run-up” between the melting and boiling points of the element was useful. An interval of 800 degrees Celsius allows, for example, to fill energy circuits with alkali metal nuclear submarines. Sodium takes heat from the reactor without boiling.

It is possible to keep the temperature within reasonable limits due to the circulation of liquid metal between the reactor and the steam generator. In the latter happens sodium cooling, water evaporates. So it turns out to rotate the turbine with the energy of high-pressure steam.

The pure element is also useful in metallurgy. What role does it play in it? sodium? Instructions The application is as follows: the metal strengthens lead-based alloys. About 1.5% of the 11th element is sufficient. Sodium is also added to alloys of other metals. It is no longer strength that is achieved, but the refractoriness of the mixtures.

Wires are worth mentioning among the products. Buy sodium for their production is more profitable than the usual one. The 11th element conducts electricity 3 times worse. But alkali metal is 9 times lighter.

This argument forced industrialists to switch to sodium buses for high currents. Thin wires continued to be made from the usual copper and.

Now, about the role sodium compounds. Peroxide not only ignites substances, but can also bleach fabrics. The hydroxide of the 11th metal is needed by oil workers. The compound purifies liquid processing products. Hydroxide is also purchased for the production of solid detergents. Without a caustic substance it is impossible to saponify the fats in them.

In parallel, the odorless white powder is also used in the production of textiles. Here the bet is on sodium chloride solution. The product can lighten both writing sheets and fabrics. The reagent is popularly called bleach.

Sodium tetraborate has become a cure for candidiasis and other vaginal infections. Sulfacyp sodium– a drug that helps with conjunctivitis and other inflammatory processes in the eyes. An antidote for salt poisoning and general intoxication of the body is sodium thiosulfate.

The food is popularly known as a remedy for heartburn. Sodium bicarbonate– an alkali that neutralizes stomach acid. The 11th element is also used for constipation. Helps out sodium sulfate.

In addition to the medical field itself, the element is also valued in cosmetology. Sodium acid- nothing more than hyaluronic acid. It maintains youthful skin. Injections are usually given into nasolabial folds and wrinkles. Acid fills them. In youth, hyaluronate is produced by the body, but with age the process slows down. You have to administer the drug from outside.

Pictured are foods containing sodium

Humanity was the first to appreciate sodium salts. But, in the 21st century, they learned the charm of the cyanide form of the metal. She helps produce. The jewel is melted in sodium cyanide. True, some other composite ores also pass into liquid form.

However, it is already easier to isolate gold from a complex melt than from solid masses. At the final stage they “connect” and... precious metal extracted

Sodium extraction

If a pure element is needed, it is extracted from ordinary table salt. Its deposits are found on all continents. If there are not enough resources, the waters of the seas are a storehouse of salt. Select from sodium chloride succeeds by melting it and subsequent electrolysis. Cathodes are made of iron or copper. Anodes are purchased from.

Sodium fluoride and potassium chloride are added to the salt. They reduce the softening temperature of raw materials from 800 to 500 degrees Celsius. This reduces sodium loss through evaporation. The method is doubly beneficial, because simultaneously with the 11th element, pure chlorine is also extracted.

Sodium price

The cost of the item depends on what is purchased sodium hydroxide solution, pure metal, its sulfate or other compound. The price tag also depends on whether it is purchased, for example, as part of a drug, or separately. There are many areas of use of sodium, hundreds and thousands of manufacturers.

Everyone has their own needs. Some substances are easily obtained in almost ready-made form, such as salt. It costs about 10-20 rubles per kilogram. Other sodium compounds need to be synthesized, which increases their cost.

One way or another, humanity is ready to pay for the benefits brought to it by the 11th element. Its mining is active and, apparently, is not going to subside.

Sodium
Atomic number	11
Appearance simple substance	silver-white soft metal
Properties of the atom
Atomic mass (molar mass)	22.989768 a. e.m. (/mol)
Atomic radius	190 pm
Ionization energy (first electron)	495.6(5.14) kJ/mol (eV)
Electronic configuration	3s 1
Chemical properties
Covalent radius	154 pm
Ion radius	97 (+1e) pm
Electronegativity (according to Pauling)	0,93
Electrode potential	-2.71 V
Oxidation states	1
Thermodynamic properties of a simple substance
Density	0.971 /cm³
Molar heat capacity	28.23 J/(mol)
Thermal conductivity	142.0 W/( ·)
Melting point	370,96
Heat of Melting	2.64 kJ/mol
Boiling point	1156,1
Heat of vaporization	97.9 kJ/mol
Molar volume	23.7 cm³/mol
Crystal lattice of a simple substance
Lattice structure	cubic body-centered
Lattice parameters	4,230
c/a ratio	—
Debye temperature	150 K

Na	11
22,98977
3s 1
Sodium

Sodium —element main subgroup of the first group, third period periodic table chemical elements D.I. Mendeleev, with atomic number 11. Denoted by the symbol Na (lat. Natrium). The simple substance sodium (CAS number: 7440-23-5) is a soft alkali metal with a silvery-white color.

In water, sodium behaves almost the same as lithium: the reaction proceeds with the rapid release of hydrogen, and sodium hydroxide is formed in the solution.

History and origin of the name

Sodium (or rather, its compounds) has been used since ancient times. For example, soda (natron), found naturally in the waters of soda lakes in Egypt. The ancient Egyptians used natural soda for embalming, bleaching canvas, cooking food, and making paints and glazes. Pliny the Elder writes that in the Nile Delta soda (it contained a sufficient proportion of impurities) was isolated from river water. It went on sale in the form of large pieces, colored gray or even black due to the admixture of coal.

Sodium was first obtained by the English chemist Humphry Davy in 1807 by electrolysis of solid NaOH.

The name "sodium" comes from the Arabic natrun in Greek - nitron and originally it referred to natural soda. The element itself was previously called Sodium.

Receipt

The first way to produce sodium was the reduction reaction sodium carbonate coal when heating a close mixture of these substances in an iron container to 1000°C:

Na 2 CO 3 +2C=2Na+3CO

Then another method of producing sodium appeared - electrolysis of molten sodium hydroxide or sodium chloride.

Physical properties

Metallic sodium stored in kerosene

Qualitative determination of sodium using a flame - bright yellow color of the emission spectrum of the “sodium D-line”, doublet 588.9950 and 589.5924 nm.

Sodium is a silvery-white metal, in thin layers with a violet tint, plastic, even soft (easily cut with a knife), a fresh cut of sodium is shiny. The electrical and thermal conductivity values of sodium are quite high, the density is 0.96842 g/cm³ (at 19.7° C), the melting point is 97.86° C, and the boiling point is 883.15° C.

Chemical properties

An alkali metal that oxidizes easily in air. To protect against atmospheric oxygen, metallic sodium is stored under a layer kerosene. Sodium is less active than lithium, therefore with nitrogen reacts only when heated:

2Na + 3N 2 = 2NaN 3

When there is a large excess of oxygen, sodium peroxide is formed

2Na + O 2 = Na 2 O 2

Application

Sodium metal is widely used in preparative chemistry and industry as a strong reducing agent, including in metallurgy. Sodium is used in the production of highly energy-intensive sodium-sulfur batteries. It is also used in truck exhaust valves as a heat sink. Occasionally, sodium metal is used as a material for electrical wires, designed for very high currents.

In an alloy with potassium, as well as with rubidium and cesium used as a highly efficient coolant. In particular, the alloy composition is sodium 12%, potassium 47 %, cesium 41% have a record low temperature melting point −78 °C and was proposed as a working fluid for ion rocket engines and coolant for nuclear power plants.

Sodium is also used in high and low pressure discharge lamps (HPLD and LPLD). NLVD lamps of the DNaT (Arc Sodium Tubular) type are very widely used in street lighting. They give off a bright yellow light. The service life of HPS lamps is 12-24 thousand hours. Therefore, gas-discharge lamps of the HPS type are indispensable for urban, architectural and industrial lighting. There are also lamps DNaS, DNaMT (Arc Sodium Matte), DNaZ (Arc Sodium Mirror) and DNaTBR (Arc Sodium Tubular Without Mercury).

Sodium metal is used in the qualitative analysis of organic matter. The alloy of sodium and the test substance is neutralized ethanol, add a few milliliters of distilled water and divide into 3 parts, J. Lassaigne's test (1843), aimed at determining nitrogen, sulfur and halogens (Beilstein test)

— Sodium chloride (table salt) is the oldest used flavoring and preservative.
— Sodium azide (Na 3 N) is used as a nitriding agent in metallurgy and in the production of lead azide.
— Sodium cyanide (NaCN) is used in the hydrometallurgical method of leaching gold from rocks, as well as in the nitrocarburization of steel and in electroplating (silvering, gilding).
— Sodium chlorate (NaClO 3) is used to destroy unwanted vegetation on railway tracks.

Biological role

In the body, sodium is found mostly outside the cells (about 15 times more than in the cytoplasm). This difference is maintained by the sodium-potassium pump, which pumps out sodium trapped inside the cell.

Together withpotassiumsodium performs the following functions:
Creating conditions for the occurrence of membrane potential and muscle contractions.
Maintaining blood osmotic concentration.
Maintaining acid-base balance.
Normalization of water balance.
Ensuring membrane transport.
Activation of many enzymes.

Sodium is found in almost all foods, although the body gets most of it from table salt. Absorption mainly occurs in the stomach and small intestine. Vitamin D improves the absorption of sodium, however, excessively salty foods and foods rich in protein interfere with normal absorption. The amount of sodium taken in from food shows the sodium content in the urine. Sodium-rich foods are characterized by accelerated excretion.

Sodium deficiency in the dieter balanced food does not occur in humans, however, some problems can arise with vegetarian diets. Temporary deficiency may be caused by diuretic use, diarrhea, excessive sweating, or excess water intake. Symptoms of sodium deficiency include weight loss, vomiting, gas in the gastrointestinal tract, and impaired absorption amino acids and monosaccharides. Long-term deficiency causes muscle cramps and neuralgia.

Excess sodium causes swelling of the legs and face, as well as increased excretion of potassium in the urine. Maximum quantity salt that can be processed by the kidneys is approximately 20-30 grams, more already life-threatening.

Sodium compounds

Sodium, Natrium, Na (11)
The name sodium - sodium, natrium comes from an ancient word common in Egypt, among the ancient Greeks (vixpov) and Romans. It is found in Pliny (Nitron) and other ancient authors and corresponds to the Hebrew neter. In ancient Egypt, natron, or nitron, was generally called an alkali obtained not only from natural soda lakes, but also from plant ash. It was used for washing, making glazes, and mummifying corpses. In the Middle Ages, the name nitron (nitron, natron, nataron), as well as boron (baurach), also applied to saltpeter (Nitrum). Arab alchemists called alkali alkali. With the discovery of gunpowder in Europe, saltpeter (Sal Petrae) began to be strictly distinguished from alkalis, and in the 17th century. already distinguished between non-volatile, or fixed alkalis, and volatile alkali (Alkali volatile). At the same time, a difference was established between vegetable (Alkali fixum vegetabile - potash) and mineral alkali (Alkali fixum minerale - soda).

At the end of the 18th century. Klaproth introduced the name Natron, or soda, for the mineral alkali, and for the vegetable alkali, Kali. Lavoisier did not place alkali in the “Table of Simple Bodies,” indicating in a note to it that these were probably complex substances that once Someday they will be decomposed. Indeed, in 1807 Davy, by electrolysis of slightly moistened solid alkalis, obtained free metals - potassium and sodium, calling them potassium and sodium. IN next year Gilbert, publisher of the famous Annals of Physics, proposed calling the new metals potassium and sodium (Natronium); Berzelius shortened the latter name to “sodium” (Natrium). At the beginning of the 19th century. in Russia sodium was called sodia (Dvigubsky, 182i; Solovyov, 1824); Strakhov proposed the name sod (1825). Sodium salts were called, for example, soda sulfate, hydrochloric soda, and at the same time acetic soda (Dvigubsky, 1828). Hess, following the example of Berzelius, introduced the name sodium.