Topic: THE SECOND GROUP OF THE PERIODIC SYSTEM
1 Feature. Atoms of elements of the 2nd group periodic system in the outer electron layer they have 2 electrons removed at a considerable distance from the nucleus. Therefore, these 2 electrons are relatively easily split off from atoms, which in this case turn into positive doubly charged ions.
The difference in the structure of the second outer layer for a number of elements of the second group determines the existence of two subgroups: the main one, including alkaline earth metals (beryllium, magnesium, calcium, strontium, barium, radium) and a secondary * subgroup, including elements: zinc, cadmium and mercury.
All elements included in the main subgroup, except for beryllium and radium, have pronounced metallic properties.
The more atomic mass, the more electropositive the metal is. So, barium is the same strong reducing agent as the alkali metals. With water, oxides of alkaline earth metals form hydroxides, the solubility of which increases from beryllium hydroxide to barium hydroxide. In the same sequence, the basic character of these compounds also increases.
The elements of the side subgroup (Zn, Cd, Hg), as well as the elements of the main subgroup, exhibit an oxidation state equal to +2, but there is also a difference between them due to different sizes of the radii of their atoms and ionization potentials.
The metallic properties of the elements of the secondary subgroup weaken from zinc to mercury. Their hydroxides are insoluble in water and have weakly basic properties.
For medicine, the elements Mg, Ca, Ba, Zn and Hg are of interest. All these elements are part of the structure of the most important drugs.
The most toxic of the elements of group II is barium in its soluble compounds, which are used only as reagents and poisons for insects and rodents. In medicine, mainly barium sulfate, a salt that is practically insoluble in any solvent, is used.
2. MAGNESIUM COMPOUNDS
Magnesium is widely distributed in nature. It does not occur in free form, but only in the form of carbonate compounds, forming minerals dolomite MgC0 3 *CaCO 3 and magnesite MgC0 3.. Magnesium is a part of talc silicates 3MgO*4Si0 2 *H 2 0, etc.
Magnesium salts are also found in the soil, in natural waters, especially sea waters, and in many mineral springs.
The value of magnesium is great. It is part of the green plant pigment - chlorophyll, participating in the process of plant photosynthesis.
Magnesium compounds play a significant role in the activity of the central nervous system of living organisms.
Physiologically, magnesium is a calcium antagonist. So, if magnesium salts cause anesthesia and paralysis, then calcium compounds remove this phenomenon. On the contrary, the action exerted by calcium compounds is removed by magnesium.
Pharmacopoeial preparations of magnesium are: magnesium oxide, burnt magnesia, basic magnesium carbonate, white magnesia, magnesium trisilicate, magnesium sulfate.
The first three drugs exhibit an antacid effect, that is, they are used for increased acidity of gastric juice. They work in the same way as very mild laxatives. Magnesium sulfate is used as a sedative, antispasmodic and laxative.
Magnesium sulfate Magnesii sulfas
MgS0 4- 7H 2 0 M. m. 246.50
Magnesium sulfate as a remedy was first used in England - English or bitter salt.
A) receiving. Magnesium sulfate is common in nature in the form of kieserite MgS0 4*7H2 0. V in large numbers magnesium sulfate is found in sea water.
Get the drug from magnesite MgC0 3 by treating it with sulfuric acid.
MgCO 3 + H 2 S0 4 > MgS0 4 + C0 2 + H 2 O
The resulting solution is concentrated by evaporation until crystallization, this gives MgS0 4 *7H 2 0.
B) Properties. Colorless prismatic crystals, weathering in the air, salty-bitter taste, odorless. It is highly soluble in water, practically insoluble in alcohol.
B) Authenticity
GF - on Mg 2+ , the formation of a precipitate of double ammonium and magnesium phosphate at
the interaction of the drug with disubstituted sodium phosphate in an ammonia solution in the presence of ammonium chloride.
MgS0 4 + Na Н P0 4 + NH 4 OH \u003d MgNH 4 P0 4 + Na 2 S0 4 + H 2 0
White
If this reaction is carried out by the drop method on a glass slide, crystals of a certain shape are formed, which can serve as confirmation of the authenticity of the drug.
GF - With organic solution 8-hydroxyquinoline, in the presence of an ammonia solution with the addition of ammonium chloride NH 4 C1, magnesium oxyquinolate is formed, which is colored greenish yellow.
GF - The sulfate ion opens with a solution of barium chloride - a white milky precipitate of barium sulfate precipitates. Insoluble in acids and alkalis.
MgS 0 4 + ВаС1 2 = Mg С1 2 + BaS 0 4
D) Purity . Arsenic not more than 0.0002%, chlorides, heavy metals, moisture are allowed.
The preparation used for injection Solutio Magnesii sulfatis 20% aut 25% pro injectionibus is tested for manganese.
GF complexometric titration method. Ammonia buffer solution is added to the drug solution, the indicator is acid chromium black special, titrated with Trilon B until the color changes from red to blue. D.b. 99%-102%
E) Application. Myotropic antispasmodic, laxative.
It is used as a laxative, 1530 g per oral intake.
When administered parenterally, magnesium sulfate has a calming effect on the central nervous system.
As an antispasmodic for hypertension in the form of a 25% solution (subcutaneously);
For labor pain relief, 1020 ml of a 25% solution intramuscularly;
As an anticonvulsant in the same doses as in labor pain relief;
As a choleretic agent 2025% solution (inside).
In case of respiratory depression associated with an overdose (kurepodnoe), a 10% solution of calcium chloride is used intravenously.
Issue: in powder, 10%, 20%, 25% solution in ampoules, 2.5, 10 and 20 ml each.
Powder in sachets 10.0-50.0. Cormagnesin, 32% Magnesium-Diasporal forte
g) Storage: dry, cool place.
3. CALCIUM COMPOUNDS
Calcium is widely distributed in nature. Due to its high chemical activity, it is found in nature only in a bound state. It occurs as numerous deposits of limestone, chalk and marble - these are natural varieties of calcium carbonate CaCO3. There are also gypsum CaS0 4 -2H 2 0, phosphorite Ca 3 (P0 4) 2 and silicates.
All natural calcium compounds, especially carbonates, serve as sources of medical calcium preparations; marble is more often used as the purest.
Calcium plays an important role in the life of the body. It is part of the dental tissue, bones, nervous tissue, muscles, blood. Calcium ions enhance the vital activity of cells, contribute to the contraction of skeletal muscles and heart muscles, are necessary for the formation bone tissue and the process of blood clotting.
With a decrease in the concentration of calcium ions in the blood, muscle excitability increases, which often leads to convulsions. Solutions of calcium salts relieve itching caused by an allergic condition, so they are classified as antiallergic drugs.
Calcium oxide (burnt lime), burnt calcium sulfate (burnt gypsum), precipitated calcium carbonate (precipitated chalk), calcium chloride and salts of organic acids (calcium glycerophosphate, calcium gluconate, etc.) are used in medicine from calcium compounds. The pharmacopoeial drug is calcium chloride.
Calcium chloride Calcii chloridum
CaCl 2 -6H 2 0 M. m. 219.08
A) receiving. Calcium chloride, intended for medical purposes, is obtained by treating calcium carbonate (marble) with hydrochloric acid.
CaCO 3 + 2HC1 \u003d CaCl 2 + C0 2 + H 2 O
Pure calcium chloride CaCl crystallizes out of water. 2 -6H 2 0.
B) Properties. It is a colorless odorless prismatic crystals, bitter-salty taste; very easily soluble in water, causing a strong cooling of the solution. Easily soluble in 95% alcohol.
The drug is very hygroscopic and deliquesces in air. At a temperature of 94°C, it melts in its water of crystallization. Aqueous solutions are neutral. When heated to 200°C, it loses some of its water of crystallization and turns into calcium chloride dihydrate CaCl 2 -2H 2 0, The hygroscopicity of the drug and its ability to spread under the influence of moisture make the composition of this salt unstable, which can lead to inaccurate dosage in the manufacture of drugs with calcium chloride. Given this, pharmacies prepare a 50% solution of it (Calcium chloratum solutum 50%) and the necessary drugs are prepared from this concentrate.
C) Authenticity:
GF - reaction to Ca 2+ reaction with ammonium oxalate,
(NH 4) 2 C 2 0 4 + CaC 1 2 \u003d CaC 2 0 4 + 2NH 4 Cl
white sediment
The precipitate is soluble in mineral acids and insoluble in acetic acid.
The formation of a white precipitate when the drug interacts with sulfuric acid or alkali metal sulfates.
CaCl 2 + H 2 S0 4 = CaS0 4 + 2HC1
white sediment
The calcium sulfate precipitate dissolves in ammonium sulfate to form a colorless complex.
GF- calcium salts color the burner flame brick red.
GF for chlorides with silver nitrate
CaCl 2 + Ag N O 3 \u003d Ag Cl + Ca (N O 3) 2
White curd sediment
D) Purity . Impurities of soluble salts of barium, iron, aluminum, phosphates are not allowed in the preparation.
Sulfates, heavy metals, arsenic, magnesium salts are allowed according to standards.
D) quantitative definition
GF - is determined complexometrically with the indicator acid chromium dark blue. When titrated with Trilon B, adding an ammonia buffer solution, the color of the solution changes from cherry red to bluish-lilac (indic eriochrome black special T). Must be at least 98.0%.
Photometric, - argentometry (Mora)
Fluorometric, - refractometry
Weight (oxalate).
E) Application. Antiallergic
As a hemostatic agent for pulmonary, gastrointestinal, nasal and uterine bleeding;
In surgical practice to increase blood clotting;
In allergic diseases (bronchial asthma, urticaria) to relieve itching;
As an antidote for poisoning with magnesium salts.
Anti-inflammatory, for colds
The drug is administered orally as a 510% solution, intravenously as a 10% solution. You can not enter subcutaneously and intramuscularly, as in this case, necrosis may occur.
Release form: powder, 10% solution in ampoules.
g) Storage. In small, well-corked glass jars with paraffin-filled stoppers, in a dry place.
4. ZINC COMPOUNDS
In nature, zinc occurs in the form of minerals: gallite ZnCO 2 and zinc blende ZnS. Zinc is found in the muscular, dental and nervous tissues of the human body. The use of zinc compounds in medicine is based on the fact that zinc gives compounds with proteins - albuminates, soluble albuminates have an effect from slightly astringent to cauterizing. Insoluble albuminates usually form a film on the tissue surface and thus promote tissue healing (drying effect).
Zinc compounds in large doses are toxic, and when applied topically, they can be used as astringents and caustic agents. When administered orally, zinc compounds cause vomiting.
Pharmacopoeial preparations of zinc are zinc oxide and zinc sulfate.
Zinc sulfate Zinci sulfas
ZnSO4 *7H 2 0 M. m. 287.54
Zinc sulfate has been used in medicine since ancient times under the name of white vitriol, in contrast to the colored copper and iron sulfate.
A) receiving. From natural ore - zinc blende ZnS, which is subjected to roasting. In this case, zinc sulfide is converted to oxide, which is then treated with dilute sulfuric acid, resulting in the formation of zinc sulfate in solution. 2 ZnS + ZO 2 \u003d 2 ZnO + 2 SO 2
ZnO + Ha 2 S 0 4 = ZnS 0 4 + 4 H 2 O
The solution containing zinc sulfate is evaporated until the salt crystallizes in the form of a heptahydrate (ZnS0 4 -7H 2 0).
B) Properties. Colorless transparent crystals or fine crystalline powder, with an astringent metallic taste, odorless, very easily soluble in water, slowly in glycerol, insoluble in alcohol. It vanishes in the air.
B) Authenticity.
GF - Sulfate ion is determined by the formation of a white precipitate.
ZnS0 4 + Ba Cl 2 = Ba S0 4 + Zn Cl 2
White milky, insoluble in acids and alkalis
GF- on Zn 2+ reaction with sodium sulfide solution produces zinc sulfide ZnS white color(unlike other salts of heavy metals).
ZnS0 4 + Na 2 S \u003d ZnS 4 + Na 2 S0 4
white sediment
GF - Zn 2+ reaction with a solution of potassium ferrocyanide a white-yellowish crystalline precipitate of a double salt precipitates, insoluble in acids, but soluble in alkalis. 3 ZnS 0 4 + 2 K 2 [Fe (CN) 6] = K 2 Zn 3 [Fe (CN) 6] 2 + 3 K 2 SO 4
Bel-yellowish
A specific reaction to zinc is the Rinman green formation reaction. ZnS 0 4 dripped onto filter paper and cobalt nitrate on top, calcined, and a characteristic green color is obtainedRinman green: CoZnO 2
With dithizone ions Zn2+ in an alkaline medium they form a red color.
D) Purity . Not d.b. impurities of iron, copper, aluminum, magnesium, calcium and other heavy metals.
Arsenic admixture allowed
E) Quantification
GF complexometry. In the presence of an ammonia buffer solution and an acidic chromium black special indicator (or eriochrome black T). Titrate with Trilon B until the color of the solution changes from cherry red to bluish-purple.
E) Application externally as an antiseptic and astringent
In eye practice in the form of 0.1; 0.25; 0.5% solutions. In eye drops, zinc sulfate is often prescribed along with boric acid.
In gynecological practice for douching in the form of a 0.10.5% solution.
At skin diseases: pimples, blackheads, dermatoses.
Rarely administered orally as an emetic.
Release form: powder, eye drops 0.1; 0.25; 0.5%, drops of zinc sulfate with boric acid. Combined: Zinkin, Zincteral
g) Storage. With caution in well-sealed jars. List B.
Zinc oxide Zinci oxydum
It is a white amorphous powder with a yellowish tint, which easily absorbs carbon dioxide from the air. A characteristic property of zinc oxide is that when calcined, it becomes yellow, and when cooled, it becomes white.
Zinc oxide is used externally in the form of powders, ointments, linements as an astringent, drying and disinfectant for skin diseases: dermatitis, prickly heat, bedsores, diaper rash, ulcers, wounds, burns.
5. MERCURY COMPOUNDS
Mercury is a liquid metal. The distribution of mercury in nature is low. It occurs in its native form, interspersed in rocks, but mainly in the form of mercury sulfide HgS (cinnabar) bright red.
Pharmacopoeial preparations are mercury compounds with an oxidation state of +2: mercury yellow oxide, mercury dichloride, mercury amidochloride, mercury oxycyanide and mercury cyanide.
Inorganic mercury preparations are used as antiseptic, diuretic and laxatives.
The antiseptic action of mercury compounds is based on the ability of the mercury ion to precipitate proteins. The diuretic action of some mercury salts is associated with
the fact that, being excreted through the kidneys, they irritate the renal epithelium and promote urination.
Similarly, mercury compounds released through the intestines and irritating it, exhibit a laxative effect.
Soluble mercury salts are highly toxic and are listed as A.
Mercury oxide yellow Hydrargyri oxydum flavum
HgO M. m. 216.59
A) receiving . Use the reaction of its precipitation from soluble salts of mercury. For this purpose, dichloride or mercury nitrate is more often used. A concentrated solution of mercury (II) salt is slowly poured into a dilute alkali solution.
Hg(NO 3 ) 2 + 2NaOH = 2NaNO 3 + HgO + H2O
bright yellow precipitate
After settling the precipitate of mercury oxide, the liquid is drained, the precipitate is washed until there is no alkaline reaction and dried. All operations should be carried out in the dark, otherwise mercuric oxide Hg may form 2 0 black.
B) Properties. Heavy fine powder of yellow or orange-yellow color. It is insoluble in water, but easily soluble in hydrochloric, nitric and acetic acids. The light gradually darkens.
C) Identity for Hg2+.
To do this, it is treated with dilute hydrochloric acid to obtain a soluble salt of mercury (II), in which the Hg cation is determined. 2+
HgO + 2HC1 \u003d HgCl 2 + H.0
HF - reaction with alkali solutions, a precipitate of yellow mercury oxide precipitates.
HgCl 2 + 2KOH > HgO + 2KS + H 2 0
bright yellow precipitate
GF - reaction with a solution of potassium iodide; a bright red precipitate of mercury diiodide is formed, which dissolves in an excess of potassium iodide.
HgCl 2 + 2Kl \u003d HgJ 2 + 2KCl HgJ 2 + 2KI > K 2
bright red colorless solution
A solution of this complex salt is known as Nessler's reagent and is used as a very sensitive reagent for NH 4+;
GF - reaction with sodium sulfide solution; a brownish-black precipitate is formed, which is insoluble in dilute nitric acid.
HgCl 2 + NaS \u003d HgS + 2NaCl
brownish-black precipitate
D) Quantitative content
GF - neutralization indirectly through interaction with potassium iodide. Under the action of yellow oxide on mercury with a solution of potassium iodide, a soluble complex salt and alkali are formed, which is titrated with acid against methyl orange. HgO + 4 KI + H 2 O > K 2 [Hgl 4] + 2KOH
KOH + HC1 \u003d KS1 + H 2 0
Rodanometric method: yellow mercury oxide is dissolved in nitric acid, and the resulting salt is titrated with ammonium thiocyanate in the presence of ferroammonium alum until a red color is obtained.
G) Application as a gentle antiseptic for the preparation of eye ointments 2%.
E) store should be used with caution in well-corked dark glass jars, since mercuric oxide may form in the light, which is detected by the darkening of the drug. List B.
Topic FIRST GROUP OF THE PERIODIC SYSTEM
1.Characteristic.All elements that make up the first group of the periodic system have only the 1st electron on the outer electron layer, which they easily give away, turning into singly charged positive ions. This explains their very high reactivity with respect to electronegative elements such as halogens.
The main subgroup includes lithium, sodium, potassium, rubidium, cesium and francium. The secondary group consists of copper, silver and gold.
The elements of the main subgroup are called alkali metals, since their oxides, when interacting with water, form strong alkalis. Alkali metal salts are used in medicine.
The most widely used in medicine are sodium and potassium salts, described above in preparations derived from halogens.
2. COMPOUNDS OF COPPER AND SILVER
A secondary subgroup of elements of the first group is copper, silver and gold. They have a tendency to complex formation, especially in copper, as well as the ability to recover from compounds to free metal, while silver is more easily reduced than copper.
Of the inorganic copper compounds, copper sulfate is used in medicine. When ingested, it has an emetic effect; as an external agent, it is used for catarrh of the mucous membranes and ulcers due to astringent, irritating and cauterizing action.
Silver belongs to the "noble" metals. In nature, it occurs mainly in the form of compounds with sulfur (Ag 2S).
The use of silver preparations in medicine is based on its bactericidal properties. It has been proven that silver ions kill gram-positive and gram-negative microorganisms, as well as viruses. Silver preparations are used in medicine internally and externally as astringent, antiseptic and cauterizing agents in the treatment of skin, urological and eye diseases.
Of the silver compounds, silver nitrate (AgNO3) has received the greatest use as a good astringent and caustic agent. In medicine, colloid preparations are also used, where silver is associated with a protein and is only partially ionized. In colloidal preparations of silver, only the disinfecting properties of silver are preserved and its cauterizing effect disappears.
All soluble compounds of copper and silver are poisonous.
3. Silver nitrate Argenti nitras
AgN0 3
A) receiving by dissolving the copper-silver alloy in nitric acid when heated. To purify the resulting silver nitrate from impurities, it is precipitated with hydrochloric acid in the form of silver chloride. The latter is reduced with zinc, and the silver, freed from impurities, is again dissolved in nitric acid.
The resulting silver nitrate is treated with a small amount of water, crystals crystallize upon standing. The separated crystals are filtered off, washed with water and dried in the dark.
B) Properties colorless transparent crystals in the form of plates or cylindrical sticks of a radiant-crystalline structure in a break. Easily soluble in water, difficult in alcohol. The crystals darken in the light.
B) Authenticity
GF - Ag+ : with hydrochloric acid or its salts, a white precipitate of silver chloride precipitates, insoluble in nitric acid and readily soluble in ammonia solution AgNO 3 + HCl \u003d AgCI + HNO 3
White
AgCl + 2NH 4 0H \u003d Cl + 2H 2 O
GF - Ag+ reduction to free silver (silver mirror formation reaction). Formaldehyde solution is added to the ammonia solution of silver oxide and the liquid is heated. After some time, a coating of metallic silver in the form of a mirror forms on the walls of the vessel.
[ Ag (NH 3 ) 2 ] OH + HCO = 2Ag + HCOOH + 4 NH 3 + 2 H 2 O
Black sediment
Ag+ with potassium chromate, and a brownish-red precipitate of silver chromate precipitates. 2AgNO 3 + K 2 Cr0 4 = AgCr0 4 + 2KNO 3
Brownish red precipitate
The precipitate is soluble in nitric acid, ammonium hydroxide, hardly soluble in acetic acid.
GF - Nitrate-ion determined with diphenylamine in con. Sulfuric acid produces a blue color
The formation of a brown ring during the interaction of silver nitrate with iron sulfate in concentrated sulfuric acid.
Nitrate ion does not decolorize potassium permanganate in an acidic environment, unlike nitrite.
D) Purity acceptable acidity limit
Heavy metal salts (lead, copper, bismuth) are not allowed.
D) quantitativecontent - by precipitation according to Folhard, titrated with ammonium thiocyanate (thiocyanate)
AgNO 3 + NH 4 SCN \u003d AgSCN + NH 4 NO,
white sediment
3NH 4 SCN + (NH 4 )Fe(S0 4 )= Fe(SCN) 3 + 2(NH 4 ) 2 S0 4
Indicator iron ammonium alum until red coloration. Should be less than 99.75%.
G) Application antiseptic and cauterizing. The latter is due to the ability of silver nitrate to coagulate proteins, turning them into insoluble compounds, which is used to cauterize wounds and ulcers. For this purpose, silver nitrate is used in the form of sticks (Stilus Argenti nitrici).
In small concentrations, it has an astringent and anti-inflammatory effect. Applied externally for erosions, ulcers, acute conjunctivitis, trachoma in the form of 2510% aqueous solutions, as well as ointments (12%). Inside is prescribed in the form of a 0.050.06% solution for gastric ulcer, chronic gastritis. Release form: powder, lapis sticks.
WFD inside 0.03 g, VSD 0.1
E) Storage in well-corked dark glass jars, as it can decompose in the light, which is detected by the darkening of the drug. List A.
4. Protargol Protargolum, Argentum proteinicum, Silver proteinate
A) receiving from silver nitrate and protein (casein, gelatin, egg white, peptone)
Protected colloid: contains silver oxide (7.88.3%) and albumin hydrolysis products.
B) Properties Light amorphous yellow-brown powder, odorless, slightly bitter, slightly astringent taste. Easily soluble in cold water, insoluble in alcohol.
B) Authenticity
GF- protein is determined by the appearance of the smell of burnt horn and charring of the drug when heated.
GF- the residue from combustion (it is white) is dissolved in HNO 3 and carry out reactions on Ag+ with chlorides.
- (biuret re-I) the drug is boiled with razb. HCl, a precipitate formed, it was filtered off, and NaOH and C were added to the clear filtrate. US O 4, a violet color appears (for protein).
D) Purity not d.b. impurities of silver compounds, protein decomposition products.
D) quantitativedefinition: after ashing the drug with sulfuric acid. Argentometry method, Folgard variant. D.b. 7.88.3%
G) Application
Antibacterial, anti-inflammatory agent. Applied externally in ophthalmology 1-2% solution (conjunctivitis, blenorrhea, blepharitis), urology 0.1-1% (washing Bladder), otorhinolaryngology (ears, nose), gynecology. Inside with stomach ulcers and intestinal diseases.
Release form: powder and LF in pharmacies.
E) Storage : according to list B. In well-sealed dark glass jars
5. Collargol (Collargolum, Argentum colloidal, Silver colloid)
Colloidal system with 70-75% content of highly dispersed metallic silver and protective proteins (casein and gelatin hydrolysates).
Greenish-black or bluish-black plates with a metallic sheen, soluble in water to form a colloidal solution. When treated with water, it swells and forms alkaline, negatively charged sols.
Antibacterial agent. Apply:
0.2 1% solutions for washing purulent wounds;
1 2% solutions for washing the bladder with chronic cystitis and urethritis,
25% solutions in the form of eye drops for the treatment of purulent conjunctivitis and blennorrhea.
With erysipelas, soft chancre, sometimes 15% of the ointment is prescribed.
Rarely in septic conditions intravenous administration.
Storage: according to list B. In well-sealed dark glass jars
16. Which of the gases taken with the same mass occupies the largest volume under the same conditions:
17. Determine molar mass equivalent (g/mol) of sulfur in sulfur oxide (VI):
18. What is the mass fraction (%) of the metal in the oxide if the molar mass of the equivalent of the trivalent metal is 15 g / mol:
19. What is the relative molecular weight of a gas if this gas is 2.2 times heavier than air:
20. Which of the following equations is called the Mendeleev-Clapeyron equation:
3)PV=RT
21. Indicate 3 gases that have the same density for any other gas:
1) CH 4, SO 2, Cl 2
2) C 2 H 4, CH 4, F 2
3) CO, Cl 2 , H 2
4) CO, C 2 H 4, N 2
5) N 2, CH 4, H 2
22. How many moles of oxygen are formed from 3 moles of potassium chlorate during its complete thermal decomposition:
23. What amount (mol) of FeS 2 will be required to obtain 64 g of SO 2 according to the equation:
4 FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2;
24. What mass (g) of calcium carbonate will be consumed to obtain 44.8 liters of carbon dioxide, measured at n.o.:
1) 200,0;
25. The equivalent of aluminum is:
1) aluminum atom;
2) 1/2 atom of aluminum;
3) 1/3 of an aluminum atom;
4) two aluminum atoms;
5) 1 mol of aluminum atoms.
26. The law of the constancy of the composition of substances is valid for substances:
1) with a molecular structure;
2) with a non-molecular structure;
3) with an ionic crystal lattice;
4) with an atomic crystal lattice;
5) for oxides and salts.
27. The equivalent of magnesium is:
1) magnesium atom;
2) 1/2 of a magnesium atom;
3) 1/3 of a magnesium atom;
4) two magnesium atoms;
5) 1 mol of magnesium atoms.
28. To neutralize 2.45 g of acid, 2.80 g of potassium hydroxide is consumed. Define
molar mass equivalent of acid:
1) 98 g/mol;
2) 36.5 g/mol;
3) 63 g/mol;
4) 40 g/mol;
g/mol.
Classification and nomenclature of inorganic compounds
1) Na 2 O; CaO; CO2
2)SO3; CuO; CrO3
3) Mn 2 O 7; CuO; CrO3
4)SO3; CO2; P2O5
5) Na 2 O; H2O; CO2
30. Only acid oxides series:
1) CO 2 ; SiO 2 ; MNO; CrO3
2) V 2 O 5 ; CrO 3 ; TeO 3 ; Mn2O7
3) CuO; SO2; NiO; MNO
4) CaO; P 2 O 3 ; Mn 2 O 7 ; Cr2O3
5) Na 2 O; H2O; CuO; Mn2O7
31. Cannot be used to neutralize sulfuric acid:
1) sodium bicarbonate;
2) magnesium oxide;
3) hydroxomagnesium chloride;
4) sodium hydrosulfate;
5) sodium oxide
32. To neutralize sulfuric acid, you can use:
2) Mg (OH) 2
33. Exhale with a glass tube carbon dioxide into solutions. The change will be in solution:
3) Ca(OH) 2 ;
34. By dissolving the corresponding oxide in water, you can get:
35. Under certain conditions, salt is formed in the case of:
1) N 2 O 5 + SO 3;
4) H 2 SO 4 + NH 3;
36. Can form acidic salts:
1) H 3 PO 4;
37. Can form basic salts:
2) Ba(OH) 2 ;
38. The mass of limestone required to obtain 112 kg of quicklime:
39. Reacts with water:
2) CaO;
40. Soluble in water:
3) Ba(OH) 2 ;
41. To obtain potassium phosphate, potassium hydrogen phosphate must be acted upon:
42. Acid oxide:
3) Mn 2 O 7;
43. Will directly interact in aqueous solution:
2) Cu(OH) 2 and ZnO;
3) AI 2 O 3 and HCI;
4) Rb 2 O and NaOH;
5) CaO and K 2 O.
44. All acidic salts in the group:
1) KCI, CuOHCI, NaHSO 4 ;
2) KAI(SO 4) 2 , Na, Ca(HCO 3) 2 ;
3) CuS, NaHSO 3 , Cu(HS) 2 ;
4) NaHCO 3 , Na 2 HPO 4 , NaH 2 PO 4 ;
5) AIOHCI 2 , NaHCO 3 , NaCN.
45. Does not form acid salts:
4) HPO 3 ;
46. Wrong title:
1) ferrous sulfate;
2) potassium sulfate;
3) iron (II) hydrochloride;
4) copper (I) chloride;
5) ammonium sulfate.
47. When water is split off from a monobasic acid weighing 16.0 g, formed by an element in the oxidation state +5, an oxide weighing 14.56 g will be obtained. The acid was taken:
1) nitrogen;
2) metavanadium;
3) orthophosphoric;
4) arsenic;
5) chlorine.
48. When calcining metal (III) weighing 10.8 g in air, a metal oxide weighing 20.4 g was obtained. For calcination, the following was taken:
2) aluminum AI;
3) iron Fe;
4) scandium Sc;
5) sodium Na.
49. Sign characterizing hydrochloric acid:
1) dibasic;
2) weak;
3) volatile;
4) oxygen-containing;
5) acid is an oxidizing agent.
50. Dibasic acid:
1) nitrogen;
2) salt;
3) acetic;
4) hydrocyanic;
Selenium.
51. Monobasic acid:
1) selenium;
2) phosphorous;
3) tellurium;
4) boric;
5) hydrocyanic.
52. Forms two types of acid salts:
1) sulfuric acid;
2) orthophosphoric acid;
3) metaphosphoric acid;
4) selenic acid;
5) sulfurous acid.
53. Does not form acid salts:
1) sulfuric acid;
2) phosphoric acid;
3) metaphosphoric acid;
4) selenic acid;
5) sulfurous acid.
54. Specify the cationic complex:
1) Na 3;
3) K3;
4) CI3;
5) K2.
55. Complex non-electrolyte:
1) Na 3;
2) ;
3) K3;
4) CI 3 ;
5) K2.
56. Anion complex:
1) potassium hexacyanoferrate (III);
2) tetrachlorodiammineplatinum (IV);
3) diamminesilver chloride;
57. Complex non-electrolyte:
1) potassium hexacyanoferrate (III);
2) tetrachlorodiammineplatinum (IV);
3) diamminesilver chloride;
4) tetraammine copper (II) sulfate;
5) hexaaquachromium (III) chloride.
58. Formula of hexaaquachromium (III) chloride:
1) Na 3;
2) CI
3) CI 2 ;
4) CI 3 ;
5) K 2 Cr 2 O 7 .
59. Formula of hexaaquachromium (II) chloride:
1) Na 3;
2) CI
3) CI 2 ; 3bl
4) CI 3 ;
5) K 2 Cr 2 O 7 .
60. Yellow blood salt refers to:
1) To aquacomplexes;
2) Hydrates;
3) To acid complexes;
4) To ammonia;
5) To chelates.
61. Copper sulfate refers to:
1) To aquacomplexes;
2) Hydrates;
3) To acid complexes;
4) To ammonia;
5) To chelates.
62. To obtain CaCO 3, add to a solution of Ca (HCO 3) 2:
1) Ca (OH) 2;
"The structure of matter and periodic law DI. Mendeleev”
63. In the nucleus of the most common isotope of lead 207 Pb neutrons:
2) 125
64. The maximum number of electrons at the level n = 3:
65. On an energy level with n = 4 sublevels:
66. Number of energy levels in a tungsten atom:
67. In the nucleus of an osmium atom, protons:
68. The nucleus of a krypton atom contains:
R and 44n
69. The number of electrons in a chromium ion:
70. An ion, which has 18 electrons and 16 protons in its composition, has a nuclear charge:
71. The maximum number of electrons that can occupy a 3s orbital:
72. The electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 has an atom:
73. Incorrect designations of orbitals:
3) 1p, 2d
74. Identical with the argon atom electronic configuration particle has:
1) Ca 2+
75. Electron affinity is called:
1) the energy required to detach an electron from an unexcited atom;
2) the ability of an atom of a given element to pull electron density onto itself;
3) the transition of an electron to a higher energy level;
4) release of energy when an electron is attached to an atom or ion;
5) chemical bond energy.
76. As a result of a nuclear reaction 
an isotope is formed:
77. In the hydrogen atom, the absorption of a photon with a minimum energy requires the transition of an electron:
78. The corpuscular-wave nature of an electron is characterized by the equation:
79. For the valence electron of the potassium atom, the values of quantum numbers (n, l, m l , m s):
1) 4, 1, -1, - :
2) 4, 1, +1, + : 3bm
3) 4, 0, 0, + :
4) 5, 0, +1, + :
80. The charge of the nucleus of an atom, in which the configuration of valence electrons in the ground state is …4d 2 5s 2:
81. The main quantum number n determines:
1) form electron cloud;
2) electron energy;
82. The orbital quantum number l determines:
1) the shape of the electron cloud;
2) electron energy;
3) orientation of the electron cloud in space;
4) rotation of an electron around its own axis;
5) hybridization of the electron cloud.
83. Magnetic quantum number m determines:
1) the shape of the electron cloud;
2) electron energy;
3) orientation of the electron cloud in space;
4) rotation of an electron around its own axis;
5) hybridization of the electron cloud.
84. The spin quantum number m s determines:
1) the shape of the electron cloud;
2) electron energy;
3) orientation of the electron cloud in space;
4) rotation of an electron around its own axis;
5) hybridization of the electron cloud.
85. During - decay, the nucleus of an atom of a radioactive element emits:
1) electron;
2) positron;
4) two protons;
5) two neutrons.
86. During - - decay, the nucleus of an atom of a radioactive element emits:
1) electron;
2) positron;
3) two protons and two neutrons combined into the nucleus of a helium atom;
4) two protons;
5) two neutrons.
87. During + - decay, the nucleus of an atom of a radioactive element emits:
1) electron;
2) positron;
3) two protons and two neutrons combined into the nucleus of a helium atom;
4) two protons;
5) two neutrons.
88. The atomic orbital has the smallest value of the sum (n + l):
89. Highest value sum (n + l) has an atomic orbital
90. The nitrogen atom will be more stable if three electrons are distributed on the 2p sublevel, one in each orbital. This matches the content:
2) the Pauli principle;
3) Gund rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
91. The twenty-first electron of the scandium atom is located on the 3d sublevel, and not on the 4p sublevel. This matches the content:
1) The principle of least energy;
2) the Pauli principle;
3) Gund rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
92. The nineteenth electron of the potassium atom is located on the 4s sublevel, and not on the 3d sublevel. This matches the content:
1) The principle of least energy;
2) the Pauli principle;
3) Gund rules;
4) 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
93. The only electron of the hydrogen atom in the ground state is located on the first energy level. This matches the content:
1) The principle of least energy;
2) the Pauli principle;
3) Gund rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
94. The maximum number of electrons at the second energy level of atoms of elements
is 8. This matches the content:
1) The principle of least energy;
2) the Pauli principle;
3) Gund rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
95. One of the mechanisms for the formation of a covalent bond:
1) radical;
2) exchange;
3) molecular;
4) ionic;
5) chain.
96. An example of a non-polar molecule having a polar covalent bond would be:
4) CCl4
97. Non-polar molecule:
98. In the series of molecules LiF - BeF 2 - BF 3 - CF 4 - NF 3 - OF 2 - F 2:
1) the nature of the connection does not change;
2) the ionic nature of the bond is enhanced;
3) the covalent nature of the bond weakens;
4) the covalent nature of the bond is enhanced;
5) there is no correct answer.
99. A covalent bond by a donor-acceptor mechanism is formed in a molecule:
2) CCl 4 ;
3) NH 4 C1;
4) NH3;
100. In a nitrogen molecule are formed:
1) only -connections;
2) only -connections;
3) both - and -connections;
4) single bond;
5) double bond.
101. The methane molecule has the structure:
1) flat;
2) tetrahedral;
3) pyramidal;
4) square;
102. The formation of an ionic lattice is characteristic of:
1) cesium iodide;
2) graphite;
3) naphthalene;
4) diamond;
103. Which of the following substances is characterized by the formation of an atomic lattice:
1) ammonium nitrate;
2) diamond;
4) sodium chloride;
5) sodium.
104. Chemical elements are arranged in ascending order of electronegativity in
1) Si, P, Se, Br, Cl, O;
2) Si, P, Br, Se, Cl, O;
3) P, Si, Br, Se, Cl, O;
4) Br, P, Cl, Si, Se;
5) Si, P, Se, Cl, O, Br
105. The valence orbitals of the beryllium atom in the beryllium hydride molecule ... are hybridized
106. The beryllium hydride molecule has the structure:
1) square
flat
3) tetrahedral
5) spherical.
107. The valence orbitals of the boron atom in the BF 3 molecule are hybridized according to the type:
108. Which of the molecules is the most durable?
109. Which of the indicated molecules has the largest dipole?
110. What is the spatial configuration of a molecule during sp 2 hybridization of AO:
1) linear
2) tetrahedron
3) flat square
Flat trigonal
111. A molecule has an octahedral structure if the next hybridization occurs
3) d2sp3
112. Modern theory the structure of the atom is based on the representations:
1) classical mechanics;
2) quantum mechanics;
3) Bohr's theory;
4) electrodynamics;
5) chemical kinetics.
113. Of the following characteristics of the atoms of the elements, they periodically change:
1) the charge of the nucleus of an atom
2) relative atomic mass;
3) the number of energy levels in an atom;
4) the number of electrons in the outer energy level;
5) total number electrons.
114. Within a period, an increase in the ordinal number of an element is usually accompanied by:
1) a decrease in the atomic radius and an increase in the electronegativity of the atom;
2) an increase in the atomic radius and a decrease in the electronegativity of the atom;
3) a decrease in the atomic radius and a decrease in the electronegativity of the atom
4) an increase in the atomic radius and an increase in the electronegativity of the atom
5) decrease in electronegativity.
115. An atom of which of the elements most easily gives one electron:
1) sodium, atomic number 11;
2) magnesium, serial number 12;
3) aluminum, serial number 13;
4) silicon, serial number 14;
5) sulfur, serial number 16.
116. Atoms of elements of the IA group of the periodic system of elements have the same number:
1) electrons in the outer electronic level;
2) neutrons;
3) all electrons;
4) electron shells;
5) protons.
117. Which of the following elements is named after the country:
118. Which row includes only transitional elements:
1) elements 11, 14, 22, 42;
2) elements 13, 33, 54, 83;
3) elements 24, 39, 74, 80;
4) elements 19, 32, 51, 101;
5) elements 19, 20, 21, 22.
119. An atom of which of the elements of the VA group has the maximum radius:
2) phosphorus;
3) arsenic;
4) bismuth;
5) antimony.
120. What series of elements is presented in ascending order of atomic radius:
1) O, S, Se, Te;
3) Na, Mg, AI, Si;
4) J, Br, CI, F;
5) Sc, Te, V, Cr.
121. The metallic nature of the properties of elements in the series Mg - Ca - Sr - Ba
1) decreases;
2) increases;
3) does not change;
4) decreases and then increases;
5) increases and then decreases.
122. The main properties of the hydroxides of elements of the JA group as the serial number increases
1) decrease,
2) increase,
3) remain unchanged,
4) decrease and then increase,
5) increase and then decrease.
123. Simple substances what elements have the greatest similarity of physical and chemical properties:
3) F, CI;
124. The existence of which of the following elements was predicted by D.I. Mendeleev:
3) Sc, Ga, Ge;
125. What distinguishes large periods from small ones:
1) the presence of alkali metals;
2) absence of inert gases;
3) the presence of d- and f-elements;
4) the presence of non-metals;
5) the presence of elements with metallic properties.
126. How to determine the period in which this element is located by the electronic formula of an element:
1) by the value of the principal quantum number of the external energy level;
2) by the number of valence electrons;
3) by the number of electrons in the external energy level;
4) by the number of sublevels in the external energy level;
5) by the value of the sublevel where the last valence electron is located.
127. Which element has the lowest ionization potential:
128. The chemical element of the third period forms the highest oxide of composition E 2 O 3 . How are electrons distributed in an atom of a given element?
1) 1s 2 2s 2 2p 1
2) 1s 2 2s 2 2p 6 3s 1
3) 1s 2 2s 2 2p 6 3s 2 3p 1
4) 1s 2 2s 2 2p 6 3s 2 3p 6
5) 1s 2 2s 2 2p 3
129. What chemical element forms a base with the most pronounced properties
1) calcium
3) aluminum
Potassium
5) beryllium
130. A chemical element has the following distribution of electrons over the electron layers in an atom 2.8.6. What position does it occupy in the periodic system of chemical elements of D.I. Mendeleev:
1) 6 period 6 group
Period 6 group
3) 2 period 6 group
4) 3 period 2 group
5) 2 period 8 group
131. The quantum numbers of the last electron in an atom of an element are n = 5, l = 1, m = -1, m s = -. Where is this element located in the periodic table?
1) 5th period, first group
2) 5th period, main subgroup 4th group
3) 4th period, sixth group
period, sixth group main subgroup
5) 5th period, the sixth group is a secondary subgroup.
132. The formula of the highest oxide of the chemical element EO 2 . To which group of the main subgroup of the periodic system of chemical elements D.I. Mendeleev belongs to this element?
Fourth
5) sixth.
133. From the above list of elements - Li, Na, Ag, Au, Ca, Ba - alkali metals include:
1) all metals;
2) Li, Na;
3) Li, Na, Ag, Au;
134. In the series from Li to Fr:
1) metallic properties are enhanced;
2) metallic properties decrease;
3) the atomic radius decreases;
4) the bond of valence electrons with the nucleus is enhanced;
5) activity in relation to water decreases
135. The sequence of elements does not apply to metals:
3) B, As, Te;
136. With an increase in the ordinal number of an element, the acidic properties of oxides N 2 O 3 - P 2 O 3 - As 2 O 3
Sb 2 O 3 - Bi 2 O 3
1) are amplified;
2) weaken;
3) remain unchanged;
4) increase, then weaken;
5) weaken, then intensify.
137. An ammonia molecule has the form:
1) curved;
2) linear;
3) planar;
4) pyramidal;
138. In the series C-Si-Ge-Sn-Pb, non-metallic signs of elements:
1) increase;
2) weaken;
3) do not change;
4) increase and then weaken;
5) weaken and then increase.
139. Valence orbitals at the carbon atom in the CH 4 methane molecule can be described on the basis of
ideas about type hybridization (sp; sp 2; sp 3; d 2 sp 3; dsp 2).
In this case, the methane molecule has the form:
1) linear;
2) flat;
3) tetrahedral;
5) square.
140. The valence orbitals of a silicon atom in a SiH 4 silane molecule can be described on the basis of the concept of hybridization of the type (sp; sp 2; sp 3; d 2 sp 3; dsp 2).
Therefore, the silane molecule has the form:
1) linear;
2) flat;
3) tetrahedral;
5) square.
141. What is the maximum number of covalent bonds that a nitrogen atom can form:
142. The nitrogen atom of an ammonia molecule with a hydrogen ion forms:
1) ionic bond;
2) covalent bond by exchange mechanism;
3) non-polar covalent bond;
4) covalent bond according to the donor-acceptor mechanism;
5) hydrogen bond.
143. Which statement is false:
4) An ionic bond has saturation;
144. Which statement is false:
1) A covalent bond has saturation;
2) The covalent bond has a direction;
3) Ionic bond has unsaturability;
4) Ionic bond has a direction;
5) The ionic bond is non-directional.
“Regularities chemical processes and their energy
145. What changes in temperature T and pressure P contribute to the formation of CO by the reaction C (solid) + CO 2 (g.) 2CO (g.) -119.8 kJ:
1) increase in T and increase in P;
2) increase in T and decrease in P;
3) decrease in T and increase in P;
4) decrease in T and decrease in P;
5) R increase.
146. How many times will the rate of a chemical reaction increase with an increase in temperature by 30 0 if the temperature coefficient of the rate is 2?
147. By how many degrees should the temperature be lowered so that the reaction rate decreases by 27 times if the temperature coefficient of the rate is 3?
148. How many times will the reaction rate X + 2Y \u003d Z increase with increasing concentration
Y 3 times?
149. How many times will the rate of the forward reaction increase compared to the rate of the reverse reaction in the 2NO + O 2 2NO 2 system with a 2-fold increase in pressure?
150. Specify the correct expression for the speed of the system: 2Cr+3Cl 2 = 2CrCl 3
5) v=k[A][C].
154. A catalyst speeds up a chemical reaction due to:
1) decrease in activation energy;
2) increase in activation energy;
3) decrease in the heat of reaction;
4) increase in concentration;
5) all answers are wrong.
155. The equilibrium of the reaction Fe 3 O 4 + 4CO "3Fe + 4CO 2 -43.7 kJ shifts to the left:
1) when the temperature drops;
2) when the temperature rises;
3) with increasing pressure;
4) with an increase in the concentration of the starting substances;
5) when adding a catalyst.
156. How many times will the rate of a chemical reaction increase with an increase in temperature by 30 0 if the temperature coefficient of the rate is 3?
157. By how many degrees should the temperature be raised so that the reaction rate increases by 27 times if the temperature coefficient of the rate is 3?
158. How many times does the rate of the reaction X + 2Y = Z increase with an increase in the concentration of X by 3 times?
159. How many times will the rate of the forward reaction increase compared to the rate of the reverse reaction in the 2CO + O 2 2CO 2 system with a 2-fold increase in pressure?
160. How will the rate of the gas reaction 2NO 2 \u003d N 2 O 4 increase with an increase in the concentration of NO 2 by 5 times?
161. How many times will the rate of the gas reaction 2NO + O 2 \u003d 2NO 2 decrease when the mixture of reacting gases is diluted 3 times?
162. By how many degrees should the temperature be lowered so that the reaction rate decreases by 81 times at a temperature coefficient of 3?
163. How many times will the reaction rate 2NO + O 2 \u003d 2NO 2 increase when the pressure in the system is increased by 4 times?
164. How many times will the rate of the forward reaction increase compared to the rate of the reverse reaction in the 2NO + O 2 2NO 2 system with an increase in pressure in the system by 5 times?
165. How will the reaction rate 2SO 2,g + O 2,g 2SO 3,g change with increasing concentration 1) will increase by 3 times; 2) increase by 9 times; 3) decrease by 3 times; 4) decrease by 9 times; 5) will not change. 166. How will the rate of the reaction 2O 3,g 3O 2,g change when the pressure is doubled? 1) will decrease by 2 times; 2) decrease by 8 times; 3) will increase by 4 times; 4) decrease by 4 times; 5) will increase by 2 times. 167. How will the reaction rate 2NO g + O 2,g 2NO 2,g change while reducing concentrations of NO and O 2 by 2 times? 1) will increase by 2 times; 2) will decrease by 2 times; 3) will increase by 24 times; 4) decrease by 24 times; Decrease by 8 times. 168. How will the rate of the direct reaction H 2 O, g H 2, g + O 2, g change if the pressure in the system increases 4 times? 1) will increase by 2 times; 2) will decrease by 2 times; 3) will not change; 4) will increase by 4 times; 5) will decrease by 4 times. 169. The law of mass action was discovered: 1) M.V. Lomonosov 2) G.I. Hess 3) J.W. Gibbs K. Guldberg and P. Waage 5) Van't Hoff 170. Which of the following systems is homogeneous Sodium chloride solution 2) ice water 3) saturated solution with a precipitate 4) coal and sulfur in the air 5) a mixture of gasoline and water 171. The value of the rate constant of a chemical reaction does not depend 1) on the nature of the reacting substances 2) on temperature 3) from the presence of catalysts From the concentration of substances 5) no factors 172. Activation energy is 1) the energy required to detach an electron from an atom 2) the excess energy of which molecules per 1 mole must have in order for their collision to lead to the formation of a new substance 3) ionization potential 4) the energy that is released as a result of the reaction 5) the energy that is released when an electron is attached to an atom. 173. The increase in the reaction rate with increasing temperature is usually characterized by: 1) the rate constant of a chemical reaction 2) chemical equilibrium constant Nitrogen, phosphorus, arsenic, antimony and bismuth belong to the main subgroup of group V of the periodic system. These elements, having five electrons in the outer layer of the atom, are generally characterized as non-metals. However, the ability to attach electrons is much less pronounced in them than in the corresponding elements of groups VI and VII. Due to the presence of five outer electrons, the highest positive oxidation of the elements of this subgroup is -5, and the negative one is 3. Due to the relatively lower electronegativity, the bond of the elements under consideration with hydrogen is less polar than the bond with hydrogen of the elements of groups VI and VII. Therefore, the hydrogen compounds of these elements do not split off hydrogen ions H in an aqueous solution, and thus do not have acidic properties. The physical and chemical properties of the elements of the nitrogen subgroup change with an increase in the serial number in the same sequence that was observed in the previously considered groups, but since the non-metallic properties are less pronounced than that of oxygen, and even more so of fluorine, the weakening of these properties upon transition to the following elements entails the appearance and growth of metallic properties. The latter are already noticeable in arsenic, antimony possesses approximately equally those and other properties, and in bismuth the metallic properties predominate over the non-metallic ones. DESCRIPTION OF THE ELEMENTS. NITROGEN(from Greek ázōos - lifeless, lat. Nitrogenium), N, a chemical element of Group V of the Mendeleev periodic system, atomic number 7, atomic mass 14.0067; colorless gas, odorless and tasteless. Historical reference. Nitrogen compounds - saltpeter, nitric acid, ammonia - were known long before nitrogen was obtained in a free state. In 1772, D. Rutherford, burning phosphorus and other substances in a glass bell, showed that the gas remaining after combustion, which he called "suffocating air," does not support respiration and combustion. In 1787, A. Lavoisier established that the "vital" and "suffocating" gases that make up the air are simple substances, and proposed the name "nitrogen". In 1784 G. Cavendish showed that nitrogen is part of saltpeter; this is where the Latin name nitrogen comes from (from the late Latin nitrum - saltpeter and the Greek gennao - I give birth, I produce), proposed in 1790 by J. A. Chaptal. By the beginning of the 19th century. the chemical inertness of nitrogen in the free state and its exceptional role in compounds with other elements as bound nitrogen were elucidated. Since then, the "binding" of nitrogen in the air has become one of the most important technical problems in chemistry. distribution in nature. Nitrogen is one of the most common elements on Earth, and most of it (about 4´1015 tons) is concentrated in the free state in the atmosphere. In the air, free nitrogen (in the form of N2 molecules) is 78.09% by volume (or 75.6% by mass), not counting minor impurities in the form of ammonia and oxides. The average content of nitrogen in the lithosphere is 1.9´10-3% by weight. Natural nitrogen compounds. - ammonium chloride NH4Cl and various nitrates (see Saltpeter.) Large accumulations of saltpeter are characteristic of a dry desert climate (Chile, Central Asia). For a long time, saltpeter was the main supplier of nitrogen for industry (now the industrial synthesis of ammonia from atmospheric nitrogen and hydrogen is of primary importance for nitrogen fixation). Small amounts of bound nitrogen are found in coal (1-2.5%) and oil (0.02-1.5%), as well as in the waters of rivers, seas and oceans. Nitrogen accumulates in soils (0.1%) and in living organisms (0.3%). Although the name "nitrogen" means "non-life-sustaining", it is actually an essential element for life. The protein of animals and humans contains 16 - 17% nitrogen. In the organisms of carnivorous animals, protein is formed due to the consumed protein substances that are present in the organisms of herbivorous animals and in plants. Plants synthesize protein by assimilating nitrogenous substances contained in the soil, mainly inorganic ones. Significant amounts of nitrogen enter the soil due to nitrogen-fixing microorganisms capable of converting free nitrogen from the air into nitrogen compounds. In nature, the nitrogen cycle is carried out, the main role in which is played by microorganisms - nitrifying, denitrifying, nitrogen-fixing, etc. However, as a result of plants extracting a huge amount of bound nitrogen from the soil (especially with intensive farming), the soils turn out to be depleted in nitrogen. Nitrogen deficiency is typical for agriculture in almost all countries, there is a nitrogen deficiency in animal husbandry ("protein starvation"). On soils poor in available nitrogen, plants develop poorly. Nitrogen fertilizers and protein feeding of animals are the most important means of advancing agriculture. Human economic activity disrupts the nitrogen cycle. For example, fuel combustion enriches the atmosphere with nitrogen, and fertilizer plants fix nitrogen in the air. Transportation of fertilizers and agricultural products redistributes nitrogen on the earth's surface. Nitrogen is the fourth most abundant element in the solar system (after hydrogen, helium and oxygen). Isotopes, atom, molecule. Natural nitrogen consists of two stable isotopes: 14N (99.635%) and 15N (0.365%). The 15N isotope is used in chemical and biochemical research as a labeled atom. Of the artificial radioactive isotopes of nitrogen, 13N has the longest half-life (T1 / 2 - 10.08 min), the rest are very short-lived. In the upper atmosphere, under the influence of cosmic radiation neutrons, 14N is converted into a radioactive isotope of carbon 14C. This process is also used in nuclear reactions to produce 14C. The outer electron shell of the nitrogen atom. consists of 5 electrons (one lone pair and three unpaired - configuration 2s22p3). Most often nitrogen. in compounds it is 3-covalent due to unpaired electrons (as in ammonia NH3). The presence of a lone pair of electrons can lead to the formation of another covalent bond, and nitrogen becomes 4-covalent (as in the ammonium ion NH4 +). Nitrogen oxidation states vary from +5 (in N205) to -3 (in NH3). Under normal conditions, in the free state, nitrogen forms an N2 molecule, where the N atoms are linked by three covalent bonds. The nitrogen molecule is very stable: its dissociation energy into atoms is 942.9 kJ / mol (225.2 kcal / mol), therefore, even at t about 3300 ° C, the degree of nitrogen dissociation. is only about 0.1%. Physical and chemical properties. Nitrogen is slightly lighter than air; density 1.2506 kg/m3 (at 0°C and 101325 n/m2 or 760 mm Hg), mp -209.86°C, tbp -195.8°C. A. liquefies with difficulty: its critical temperature is rather low (-147.1 ° C), and its critical pressure is high 3.39 MN / m2 (34.6 kgf / cm2); the density of liquid nitrogen is 808 kg (m3. In water, nitrogen is less soluble than oxygen: at 0 ° C, 23.3 g of nitrogen dissolves in 1 m3 of H2O. Better than in water, nitrogen is soluble in some hydrocarbons. Only with such active metals as lithium, calcium, magnesium, nitrogen interacts when heated to relatively low temperatures. Nitrogen reacts with most other elements at high temperatures and in the presence of catalysts. Compounds of nitrogen with oxygen N2O, NO, N2O3, NO2 and N2O5 are well studied. Of these, with the direct interaction of the elements (4000 ° C), NO oxide is formed, which, when cooled, is easily oxidized further to NO2 dioxide. In air, nitrogen oxides are formed during atmospheric discharges. They can also be obtained by the action of ionizing radiation on a mixture of nitrogen and oxygen. When nitrous N2O3 and nitric N2O5 anhydrides are dissolved in water, nitrous acid HNO2 and nitric acid HNO3 are obtained, respectively, which form salts - nitrites and nitrates. Nitrogen combines with hydrogen only at high temperature and in the presence of catalysts, and ammonia NH3 is formed. In addition to ammonia, numerous other compounds of nitrogen with hydrogen are also known, for example, hydrazine H2N-NH2, diimide HN-NH, nitric acid HN3(H-N-NºN), octazone N8H14, etc.; most nitrogen-hydrogen compounds have been isolated only in the form of organic derivatives. Nitrogen does not directly interact with halogens, therefore all nitrogen halides are obtained only indirectly, for example, nitrogen fluoride NF3- when fluorine interacts with ammonia. As a rule, nitrogen halides are low-resistant compounds (with the exception of NF3); more stable nitrogen oxyhalides - NOF, NOCI, NOBr, N02F and NO2CI. With sulfur, there is also no direct connection of nitrogen; nitrogenous sulfur N4S4 is obtained by the reaction of liquid sulfur with ammonia. When hot coke reacts with nitrogen, cyanogen (CN) is formed;. By heating nitrogen with acetylene C2H2 to 1500°C, hydrogen cyanide HCN can be obtained. The interaction of nitrogen with metals at high temperatures leads to the formation of nitrides (for example, Mg3N2). The nitrogen subgroup consists of five elements: nitrogen, phosphorus, arsenic, antimony and bismuth. These are the p-elements of group V of the periodic system of D. I. Mendeleev. Therefore, the highest degree of oxidation of these elements is +5, the lowest is -3, and +3 is also characteristic. Thus, phosphorus, arsenic, antimony and bismuth in an excited state have 5 unpaired electrons, and their valency in this state is 5. From nitrogen to bismuth, the atomic radii increase and the ionization potentials decrease. The reducing properties of neutral atoms increase from N to Bi, while the oxidizing properties weaken (see Table 21). With oxygen, elements of the nitrogen subgroup form oxides of the general formula R2O3 and R2O5. The oxides correspond to the acids HRO2 and HRO3 (and the orthoacids H3RO4, except for nitrogen). Within the subgroup, the nature of the oxides varies as follows: N2O3 - acidic oxide; P4O6 - weakly acidic oxide; As2O3 - amphoteric oxide with a predominance of acidic properties; Sb2O3 - amphoteric oxide with a predominance of basic properties; Bi2O3 is the main oxide. Thus, the acidic properties of oxides of the composition R2O3 and R2O5 decrease with an increase in the atomic number of the element. Nitrogen. Receipt In laboratories, it can be obtained by the decomposition reaction of ammonium nitrite: The reaction is exothermic, releasing 80 kcal (335 kJ), so cooling of the vessel is required during its course (although ammonium nitrite is required to start the reaction). In practice, this reaction is carried out by adding dropwise a saturated solution of sodium nitrite to a heated saturated solution of ammonium sulfate, while the ammonium nitrite formed as a result of the exchange reaction instantly decomposes. The gas released in this case is contaminated with ammonia, nitric oxide (I) and oxygen, from which it is purified by successively passing through solutions of sulfuric acid, iron (II) sulfate and over hot copper. The nitrogen is then dried. Another laboratory method for obtaining nitrogen is by heating a mixture of potassium dichromate and ammonium sulfate (in a ratio of 2:1 by weight). The reaction goes according to the equations: The purest nitrogen can be obtained by decomposition of metal azides: The so-called "air", or "atmospheric" nitrogen, that is, a mixture of nitrogen with noble gases, is obtained by reacting air with hot coke, and the so-called "generator" or "air" gas is formed - raw materials for chemical synthesis and fuel . If necessary, nitrogen can be separated from it by absorbing carbon monoxide. Molecular nitrogen is produced industrially by fractional distillation of liquid air. This method can also be used to obtain "atmospheric nitrogen". Nitrogen plants and stations that use the method of adsorption and membrane gas separation are also widely used. One of the laboratory methods is passing ammonia over copper (II) oxide at a temperature of ~700 °C: Ammonia is taken from its saturated solution by heating. The amount of CuO is 2 times more than the calculated one. Immediately before use, nitrogen is purified from oxygen and ammonia impurities by passing over copper and its oxide (II) (also ~700 °C), then dried with concentrated sulfuric acid and dry alkali. The process is rather slow, but worth it: the gas is very pure. Periodicity in changing the properties of chemical elements based on the electronic structure of their atoms Therefore, the methodical method of compiling electronic formulas of elements based on the periodic system consists in the fact that we sequentially consider the electron shell of each element along the path to the given one, identifying by its “coordinates” where its next electron went in the shell. The first two elements of the first period, hydrogen H and helium, do not belong to the s-family. Two of their electrons go to the s-sublevel of the first level. We write down: The first period ends here, the first energy level also. The next two elements of the second period, lithium Li and beryllium Be, are in the main subgroups of groups I and II. These are also s-elements. Their next electrons will be located on the s sublevel of the 2nd level. We write down Next, 6 elements of the 2nd period follow in a row: boron B, carbon C, nitrogen N, oxygen O, fluorine F and neon Ne. According to the location of these elements in the main subgroups of III - Vl groups, their next six electrons will be located on the p-sublevel of the 2nd level. We write down: The second period ends with the inert element neon, the second energy level is also completed. This is followed by two elements of the third period of the main subgroups of groups I and II: sodium Na and magnesium Mg. These are s-elements and their next electrons are located on the s-sublevel of the 3rd level. Then there are six elements of the 3rd period: aluminum Al, silicon Si, phosphorus P, sulfur S, chlorine C1, argon Ar. According to the location of these elements in the main subgroups of groups III - VI, their next electrons, among six, will be located on the p-sublevel of the 3rd level - The 3rd period is completed by the inert element argon, but the 3rd energy level is not yet completed, while there are no electrons on its third possible d-sublevel. This is followed by 2 elements of the 4th period of the main subgroups of groups I and II: potassium K and calcium Ca. These are again s-elements. Their next electrons will be at the s-sublevel, but already at the 4th level. It is energetically more profitable for these next electrons to start filling the 4th level, which is more distant from the nucleus, than to fill the 3d sublevel. We write down: The following ten elements of the 4th period from No. 21 scandium Sc to No. 30 zinc Zn are in side subgroups III - V - VI - VII - VIII - I - II groups. Since they are all d-elements, their next electrons are located on the d-sublevel before the outer level, i.e., the third from the nucleus. We write down: The following six elements of the 4th period: gallium Ga, germanium Ge, arsenic As, selenium Se, bromine Br, krypton Kr - are in the main subgroups III - VIIJ of groups. Their next 6 electrons are located on the p-sublevel of the outer, i.e., 4th level: 3b elements are considered; the fourth period is completed by the inert element krypton; completed and the 3rd energy level. However, at level 4, only two sublevels are completely filled: s and p (out of 4 possible). This is followed by 2 elements of the 5th period of the main subgroups of I and II groups: No. 37 rubidium Rb and No. 38 strontium Sr. These are elements of the s-family, and their next electrons are located on the s-sublevel of the 5th level: The last 2 elements - No. 39 yttrium YU No. 40 zirconium Zr - are already in side subgroups, i.e., belong to the d-family. Two of their next electrons will go to the d-sublevel, before the outer, i.e. Level 4 Summing up all the entries in sequence, we compose the electronic formula for the zirconium atom No. 40 The derived electronic formula for the zirconium atom can be somewhat modified by arranging the sublevels in the order of numbering their levels: The derived formula can, of course, be simplified into the distribution of electrons only over energy levels: Zr – 2|8| 18 |8 + 2| 2 (the arrow indicates the entry point of the next electron; valence electrons are underlined). The physical meaning of the category of subgroups lies not only in the difference in the place where the next electron enters the shell of the atom, but also in the levels at which the valence electrons are located. From a comparison of simplified electronic formulas, for example, chlorine (3rd period, main subgroup of group VII), zirconium (5th period, secondary subgroup of group IV) and uranium (7th period, lanthanide-actinide subgroup) №17, С1-2|8|7 №40, Zr - 2|8|18|8+ 2| 2 №92, U - 2|8|18 | 32 |18 + 3|8 + 1|2 it can be seen that for elements of any main subgroup, only electrons of the outer level (s and p) can be valence. For elements of secondary subgroups, electrons of the outer and partially pre-external level (s and d) can be valence. In lanthanides and especially actinides, valence electrons can be located at three levels: external, pre-external, and pre-external. As a rule, the total number of valence electrons is equal to the group number.
At the outer energy level, the atoms of these elements contain five electrons, which have the configuration ns2np3 and are distributed as follows:
The presence of three unpaired electrons at the outer level indicates that in the unexcited state, the atoms of the elements have a valence of 3. The outer level of the nitrogen atom consists of only two sublevels - 2s and 2p. The atoms of the remaining elements of this subgroup have vacant cells of the d-sublevel on the outer energy levels. Consequently, one of the s-electrons of the outer level can, upon excitation, go to the d-sublevel of the same level, which leads to the formation of 5 unpaired electrons.
outer electron shell of phosphorus (unexcited atom)
the outer electron shell of an excited phosphorus atom.
In the nitrogen atom, it is impossible to excite an electron in this way due to the absence of a d-sublevel at the second level. Therefore, nitrogen cannot be pentavalent, but it can form a fourth covalent bond by the donor-acceptor mechanism due to the lone electron pair 2s2. Another process is possible for the nitrogen atom. When one of the two 2s electrons is detached, nitrogen passes into a singly charged tetravalent N+ ion.
With hydrogen, nitrogen, phosphorus, and arsenic form polar RH3 compounds, exhibiting a negative oxidation state of -3. RH3 molecules have a pyramidal shape. In these compounds, the bonds of the elements with hydrogen are stronger than in the corresponding compounds of the elements of the oxygen subgroup and especially the halogen subgroup. Therefore, hydrogen compounds of elements of the nitrogen subgroup in aqueous solutions do not form hydrogen ions.
As can be seen from Table. 21, within the subgroup from nitrogen to bismuth, non-metallic properties decrease and metallic properties increase. In antimony, these properties are expressed in the same way, in bismuth, metallic properties predominate, and in nitrogen, non-metallic properties. Phosphorus, arsenic and antimony form several allotropic compounds.