Salts of nitrous and nitric acids
nitrogen fertilizers
Grade 9
Type of lesson - learning new material.
Type of lesson- conversation.
Goals and objectives of the lesson.
Educational. To acquaint students with the methods of obtaining, properties and applications of nitrates and nitrites. Consider the problem of high nitrate content in agricultural products. Give an idea of nitrogen fertilizers, their classification and representatives.
Educational. Continue developing skills: highlight the main thing, establish cause-and-effect relationships, take notes, conduct an experiment, apply knowledge in practice.
Educational. To continue the formation of a scientific worldview, the upbringing of a positive attitude towards knowledge.
Methods and methodological techniques. Independent work of students with popular science literature, preparation of reports, laboratory experiments and a demonstration experiment, a dialogical method of presenting knowledge with research elements, current control of knowledge using a test.
Lesson structure.
Announcement of the topic, goals.
Homework message and comments.
Presentation of new material (heuristic conversation based on experiment).
Current control of knowledge with the help of a test.
Summing up the lesson.
Equipment and reagents. Safety poster; tables "Decomposition of nitrates during heating", "Classification of nitrogen fertilizers", "Displacement series of acids"; test "Nitrogen and its compounds" (two options); task cards.
For demo experiment: a demonstration stand for test tubes, a spirit lamp, matches, a holder for test tubes, crucible tongs, an iron spoon for burning substances, a torch, an iron sheet for burning black powder, large test tubes, cotton wool soaked in a concentrated solution of alkali, a cup with sand, three laboratory stands; concentrated solutions of sodium hydroxide and sulfuric acid, crystalline salts - potassium nitrate, copper (II) nitrate, silver nitrate; charcoal, copper plate, sulfur, a solution of diphenylamine in concentrated sulfuric acid (dark bottle, 0.1 g of diphenylamine per
10 ml H 2 SO 4 (conc.); solutions of potassium iodide, dilute sulfuric acid, potassium nitrite; in demonstration test tubes - vegetable juices of cabbage, zucchini, pumpkin; starch iodine paper.
For laboratory experiments: a test tube with two zinc granules, three empty test tubes, glass rods, two test tubes with crystalline nitrates (the size of a pea) - barium nitrate and aluminum nitrate, litmus, solutions of copper (II) nitrate, silver nitrate, hydrochloric acid, barium chloride, distilled water .
Epigraph.“No science needs experiment as much as chemistry” (Michael Faraday).
DURING THE CLASSES
Safety Information
All nitrates are flammable. It is necessary to store nitrates separately from organic and not organic matter. All experiments with the formation of nitric oxide (IV) must be carried out in large test tubes closed with cotton swabs moistened with a concentrated alkali solution. Nitric acid should be stored in dark bottles, protected from fire. Nitrites are especially toxic.
Homework
O.S. Gabrielyan's textbook "Chemistry-9", § 26, exercise. 7. Strong students receive individual tasks.
Individual tasks
1. Translate the following entry from the alchemical language: ““Strong vodka” devours the “moon”, releasing a “fox tail”. The thickening of the resulting liquid generates a "hell stone" that blackens fabric, paper and hands. In order for the “moon” to rise again, bake the “hell stone” in the furnace.
Answer.
"Hellstone" - silver nitrate - decomposes when heated to form silver - "the moon has risen":
2AgNO 3 (cr.) 2Ag + 2NO 2 + O 2.
2.
An old scientific treatise describes the experience of obtaining a "red precipitate" *: "Mercury is dissolved in nitric acid, the solution is evaporated and the residue is heated until it becomes" red "". What is a "red precipitate"? Write the equations for the reactions leading to its formation, taking into account that mercury in the resulting compounds has an oxidation state of +2 and that the action of nitric acid on mercury releases a gas that turns brown in air.
Answer. Reaction equations:
Mercury(II) oxide HgO depending on the method of obtaining it is red or yellow color (Hg 2 O - black color). Mercury does not oxidize in air at room temperature. With prolonged heating, mercury combines with atmospheric oxygen, forming red mercury (II) oxide - HgO, which, when heated more strongly, decomposes again into mercury and oxygen:
2HgO \u003d 2Hg + O 2.
Learning new material
Composition and nomenclature of salts of nitric acid
Teacher. What do the Latin name "nitrogenium" and the Greek "nitrate" mean?
Student. "Nitrogenium" means "giving birth to saltpeter", and "nitrate" means "saltpeter".
Teacher. Potassium, sodium, calcium and ammonium nitrates are called saltpeters. For example, saltpeter: KNO 3 - potassium nitrate (Indian saltpeter), NaNO 3 - sodium nitrate (Chilean saltpeter), Ca(NO 3) 2 - calcium nitrate (Norwegian saltpeter), NH 4 NO 3 - ammonium nitrate (ammonium or ammonium nitrate, there are no deposits of it in nature). German industry is considered the first in the world to receive salt NH4NO3 from nitrogen N 2 air and hydrogen water suitable for plant nutrition.
Physical properties nitrates
Teacher. We learn about the relationship between the structure of a substance and its properties from laboratory experience..
Physical properties of nitrates
Exercise. Two test tubes contain crystalline nitrates: Ba(NO 3) 2 and Al(NO 3) 3. Pour 2 ml of distilled water into each test tube, mix with a glass rod. Observe the process of dissolution of salts. Solutions should be stored until the nature of the medium is studied.
Teacher. What are called salts?
Student. Salt is complex substances, consisting of metal ions and ions of acidic residues.
Teacher. It is necessary to build a logical chain: view chemical bond– type of crystal lattice – interaction forces between particles at lattice sites – physical properties of substances.
Student. Nitrates belong to the class of salts, so they are characterized by ionic bond and ionic crystal cell where ions are held together by electrostatic forces. Nitrates - solid crystalline substances, refractory, soluble in water, strong electrolytes.
Obtaining nitrates and nitrites
Teacher. Name ten ways to obtain salts based on the chemical properties of the most important classes of inorganic compounds..
Student.
1) Metal + non-metal = salt;
2) metal + acid = salt + hydrogen;
3) metal oxide + acid = salt + water;
4) metal hydroxide + acid = salt + water;
5) metal hydroxide + acid oxide = salt + water;
6) metal oxide + non-metal oxide = salt;
7) salt 1 + metal hydroxide (alkali) = salt 2 + metal hydroxide (insoluble base);
8) salt 1 + acid (strong) = salt 2 + acid (weak);
10) salt 1 + metal (active) = salt 2 + metal (less active).
Specific ways to obtain salts:
12) salt 1 + non-metal (active) = salt 2 + non-metal (less active);
13) amphoteric metal+ alkali \u003d salt + hydrogen;
14) non-metal + alkali \u003d salt + hydrogen.
A specific way to obtain nitrates and nitrites:
nitric oxide (IV) + alkali \u003d salt1 + salt2 + water, for example (writes on the board):
This is a redox reaction, its type is dismutation, or disproportionation.
In the presence of oxygen from NO 2 and NaOH it turns out not two salts, but one:
The type of redox reaction is intermolecular.
Teacher. Why should experiments with the formation of nitric oxide (IV) be carried out in large test tubes closed with cotton swabs moistened with aqueous alkali?
Student. Nitric oxide (IV) is a poisonous gas, it interacts with alkali and is rendered harmless.
Chemical properties of nitrates
Students perform laboratory experiments according to the printed method.
Properties of nitrates in common with other salts
The interaction of nitrates with metals,
acids, alkalis, salts
Exercise. Mark the signs of each reaction, write down the molecular and ionic equations, corresponding to the schemes:
Cu(NO 3) 2 + Zn ...,
AgNO 3 + HCl ...,
Cu(NO 3) 2 + NaOH ...,
AgNO 3 + BaCl 2 ....
Hydrolysis of nitrates
Exercise. Determine the reaction of the environment of the proposed solutions of salts: Ba (NO 3) 2 and Al (NO 3) 3. Write down the molecular and ionic equations of possible reactions indicating the medium of the solution.
Specific properties of nitrates and nitrites
Teacher. All nitrates are thermally unstable. When heated they decompose with the formation of oxygen. The nature of other reaction products depends on the position of the metal forming the nitrate in the electrochemical series of voltages:
Special position occupies ammonium nitrate, which decomposes without a solid residue:
NH 4 NO 3 (cr.) N 2 O + 2H 2 O.
The teacher does demonstration experiments.
Experience 1. Decomposition of potassium nitrate. Place 2–3 g of crystalline potassium nitrate in a large test tube, heat until the salt melts. Throw charcoal preheated in an iron spoon into the melt. Students watch a bright flash and burning coal. Substitute a cup with sand under the test tube.
Teacher. Why does an ember dipped in molten potassium nitrate burn instantly?
Student. Saltpeter decomposes with the formation of oxygen gas, so the preheated coal instantly burns in it:
C + O 2 \u003d CO 2.
Experiment 2. Decomposition of copper(II) nitrate. Place crystalline copper(II) nitrate (the size of a pea) into a large test tube, close the test tube with a cotton swab moistened with a concentrated alkali solution. Fix the tube in a rack horizontally and heat.
Teacher. Look for signs of a reaction.
Students observe the formation of brown gas NO 2 and black oxide of copper (II) CuO.
The student at the blackboard writes the reaction equation:
The type of redox reaction is intramolecular.
Experience 3. Decomposition of silver nitrate. Incandescent in a test tube, closed with a cotton swab moistened with a concentrated solution of alkali, a few crystals of silver nitrate.
Teacher. What gases are released? What is left in the test tube?
The student at the blackboard answers questions, draws up an equation for the reaction:
The type of redox reaction is intramolecular. A solid residue remained in the test tube - silver.
Teacher. Qualitative reaction to nitrate ion NO 3 - - interaction of nitrates with metallic copper when heated in the presence of concentrated sulfuric acid or with a solution of diphenylamine in H 2 SO 4 (conc.).
Experience 4. Qualitative reaction to the NO ion 3 - . Place a cleaned copper plate, a few crystals of potassium nitrate, and a few drops of concentrated sulfuric acid into a large dry test tube. Close the test tube with a cotton swab moistened with a concentrated alkali solution and heat.
Teacher. List the signs of a reaction.
Student. Brown vapors of nitric oxide (IV) appear in the test tube, which is better observed on a white screen, and greenish crystals of copper (II) nitrate appear at the copper-reaction mixture boundary..
Teacher(demonstrates a scheme for reducing the relative strength of acids). In accordance with a number of acids, each previous acid can displace the next one from the salt..
The student at the blackboard composes the reaction equations:
KNO 3 (cr.) + H 2 SO 4 (conc.) \u003d KHSO 4 + HNO 3,
The type of redox reaction is intermolecular.
Teacher. The second qualitative reaction to the nitrate ion NO 3 - we will spend a little later, when studying the content of nitrates in food.
Qualitative reaction to nitrite ion NO 2 -– interaction of nitrites with a solution of potassium iodide KI acidified with dilute sulfuric acid.
Experience 5. Qualitative reaction to the NO ion 2 - . Take 2-3 drops of potassium iodide solution, acidified with dilute sulfuric acid, and add a few drops of potassium nitrite solution. Nitrites in acidic environment are able to oxidize the iodide ion I - to free I 2, which is detected by starch iodine paper soaked in distilled water.
Teacher. How should starch iodine paper change color under the action of free I2?
Student. simple substance I 2 detected by blue starch.
The teacher writes the reaction equation:
Teacher. In this reaction NO 2 - is an oxidizing agent. However, there are other qualitative reactions to the ion NO 2 - in which it is a reducing agent. From this it can be concluded that the ion NO 3 - exhibits only oxidizing properties, and the ion NO 2 - - both oxidizing and reducing properties.
The use of nitrates and nitrites
Teacher(asks a challenging question). Why is there a lot of nitrogen in nature (it is part of the atmosphere), and plants often give a poor harvest due to nitrogen starvation?
Student. Plants cannot absorb molecular nitrogen N 2 from the air. This is the "bound nitrogen" problem. With a lack of nitrogen, the formation of chlorophyll is delayed, so the plants have a pale green color, as a result, the growth and development of the plant is delayed. Nitrogen is vital important element. Without protein there is no life, and without nitrogen there is no protein..
Teacher. What are the ways of assimilation of atmospheric nitrogen.
Student. Part of the bound nitrogen enters the soil during thunderstorms. The chemistry of the process is:
Teacher. What plants are able to increase soil fertility and what is their peculiarity?
Student. These plants (lupine, alfalfa, clover, peas, vetch) belong to the legume family (butterflies), on the roots of which nodule bacteria develop that can bind atmospheric nitrogen, converting it into compounds available to plants..
Teacher. When harvesting, a person annually carries away with them huge amounts of bound nitrogen. He covers this loss by introducing not only organic, but also mineral fertilizers (nitrate, ammonia, ammonium). Nitrogen fertilizers are applied to all crops. Nitrogen is taken up by plants in the form of the ammonium cation.and nitrate anion NO 3 -.
The teacher demonstrates the "Classification of nitrogen fertilizers" scheme.
Scheme
Teacher. One of the important characteristics is the nutrient content of the fertilizer. The calculation of the nutrient element for nitrogen fertilizers is carried out according to the nitrogen content.
 |
Plants that fix atmospheric nitrogen |
A task. What is the mass fraction of nitrogen in liquid ammonia and ammonium nitrate?
The formula for ammonia is NH 3.
Mass fraction of nitrogen in ammonia:
(N) = A r(N)/ M r(NH 3) 100%,
(N) = 14/17 100% = 82%.
The formula of ammonium nitrate is NH 4 NO 3.
Mass fraction of nitrogen in ammonium nitrate:
(N) = 2 A r(N)/ M r(NH 4 NO 3) 100%,
The impact of nitrates on the environment and the human body
1st student.Nitrogen as the main nutrient affects the growth of vegetative organs - green stems and leaves. Nitrogen fertilizers are not recommended to be applied in late autumn or early spring, because. melt water wash off half of the fertilizer. It is important to observe the norms and terms of fertilizer application, to apply them not immediately, but in several steps. Apply slow-acting forms of fertilizer (granules covered with a protective film), when planting, use varieties prone to low accumulation of nitrates. The utilization rate of nitrogen fertilizers is 40–60%. Excessive use of nitrogen fertilizers not only leads to the accumulation of nitrates in plants, but also leads to pollution of water bodies and groundwater. Anthropogenic sources of water pollution with nitrates are also metallurgy, chemical, including pulp and paper, and food industries. One of the signs of water pollution is the "bloom" of water caused by the rapid reproduction of blue-green algae. It occurs especially intensively during the melting of snow, summer and autumn rains. The maximum allowable concentration (MPC) of nitrates is regulated by GOST. For the sum of nitrate ions in the soil, the value of 130 mg/kg is accepted, in the water of different water sources - 45 mg/l.(Students write in notebooks: MPC (NO 3 - in soil) - 130 mg / kg, MPC (NO 3 - in water) - 45 mg / l.)
For the plants themselves, nitrates are harmless, but for humans and herbivores they are dangerous. The lethal dose of nitrates for humans is 8–15 g, the allowable daily intake is 5 mg/kg. Many plants are able to accumulate large amounts of nitrates, for example: cabbage, zucchini, parsley, dill, table beet, pumpkin, etc.
Such plants are called nitrate accumulators. 70% of nitrates enter the human body with vegetables, 20% with water, 6% with meat and fish. Once in the human body, part of the nitrates is absorbed in the gastrointestinal tract unchanged, the other part, depending on the presence of microorganisms, the pH value and other factors, can turn into more toxic nitrites, ammonia, hydroxylamine NH 2 OH ; secondary nitrosamines can be formed in the intestines from nitrates R 2 N–N=O with high mutagenic and carcinogenic activity. Signs of a slight poisoning are weakness, dizziness, nausea, indigestion, etc. Working capacity decreases, loss of consciousness is possible.
In the human body, nitrates interact with blood hemoglobin, turning it into methemoglobin, in which iron is oxidized to Fe 3+ and cannot serve as an oxygen carrier. That is why one of the signs of acute nitrate poisoning is cyanosis of the skin. A direct relationship has been revealed between the occurrence of malignant tumors and the intensity of nitrate intake into the body with their excess in the soil.
An experience. The study of the content of nitrates in food
(qualitative reaction to the nitrate ion NO 3 -)
Place 10 ml of vegetable juice of cabbage, zucchini, pumpkin (on a white background) into three large demonstration tubes. Pour a few drops of a solution of diphenylamine in concentrated sulfuric acid into each test tube.
The blue color of the solution will indicate the presence of nitrate ions:
NO 3 - + diphenylamine substance of intense blue color.
The blue color was present only in vegetable marrow juice, and the color was not intense blue. Consequently, the content of nitrates in zucchini is insignificant, and even less in cabbage with pumpkin.
First aid for nitrate poisoning
2nd student.First aid for nitrate poisoning is a copious gastric lavage, activated charcoal, saline laxatives - Glauber's salt Na 2 SO 4 10H 2 O and Epsom salts (bitter salt) MgSO 4 7H 2 O , fresh air.
It is possible to reduce the harmful effects of nitrates on the human body with the help of ascorbic acid (vitamin C); if its ratio with nitrates is 2:1, then nitrosamines are not formed. It has been proven that, first of all, vitamin C, as well as vitamins E and A, are inhibitors - substances that prevent and inhibit the conversion of nitrates and nitrites in the human body. It is necessary to introduce more black and red currants, other berries and fruits into the diet (by the way, there are practically no nitrates in hanging fruits). And another natural nitrate neutralizer in the human body is green tea..
Reasons for the accumulation of nitrates in vegetables
and methods of growing organic
crop production
3rd student. Nitrogen is absorbed most intensively during the growth and development of stems and leaves. When the seeds ripen, the consumption of nitrogen from the soil practically stops. Fruits that have reached full maturity no longer contain nitrates - there is a complete conversion of nitrogen compounds into proteins. But for many vegetables, it is the immature fruit (cucumbers, zucchini) that is valued. It is advisable to fertilize such crops with nitrogen fertilizers no later than 2-3 weeks before harvesting. In addition, the complete conversion of nitrates into proteins is hindered by poor lighting, excessive humidity and imbalance of nutrients (lack of phosphorus and potassium). You should not get carried away with off-season greenhouse vegetables. For example, 2 kg of greenhouse cucumbers eaten at one time can cause life-threatening nitrate poisoning. You also need to know in which parts of the plant nitrates accumulate: in cabbage - in the stalk, in carrots - in the core, in zucchini, cucumbers, watermelons, melons, potatoes - in the peel. Melon and watermelon should not eat the immature flesh adjacent to the rind. It is better to peel cucumbers and cut off the place where they are attached to the stem. In green crops, nitrates accumulate in the stems (parsley, lettuce, dill, celery). The content of nitrates in different parts of plants is uneven: in petioles, stems, roots, their content is 1.5–4.0 times higher than in leaves. The World Health Organization considers acceptable content nitrates in dietary products up to 300 mg NO 3 - per 1 kg of raw material.(Students write in their notebooks: MPC (NO 3 - in dietary products) - 300 mg / kg.)
If the most high content nitrates are noted in beets, cabbage, lettuce, green onions, the lowest content of nitrates is in onions, tomatoes, garlic, peppers, and beans.
In order to grow environmentally friendly products, first of all, it is necessary to correctly apply nitrogen fertilizers to the soil: in strictly calculated doses and at optimal times. It is necessary to grow vegetables, especially green crops, in good light, optimal indicators of soil moisture and temperature. And yet, to reduce the content of nitrates, it is better to feed vegetable crops with organic fertilizers. Untimely application of fertilizers, especially in excess doses, including organic fertilizer - manure, leads to the fact that the mineral nitrogen compounds that have entered the plant do not have time to completely turn into protein ones.
4th student.In spring, green crops appear on the shelves of shops and markets: lettuce, spinach, green onions, cucumbers grown in a greenhouse, in closed ground. How to reduce the content of nitrates in them? Let's list some of them.
1. Such early crops as parsley, dill, celery must be placed as a bouquet in water on a straight line. sunlight. Under such conditions, nitrates in the leaves are completely processed within 2–3 hours and then are practically not detected. After that, greens can be safely used in writing.
2. Before cooking, beets, zucchini, pumpkin must be cut into small cubes and poured 2-3 times with warm water, keeping for 5-10 minutes. Nitrates are highly soluble in water, especially warm water, and are washed out with water (see the table of solubility of acids, bases, salts). When washing and cleaning, 10-15% of nitrates are lost.
3. Boiling vegetables reduces nitrate content by 50-80%.
4. Reduces the amount of nitrates in vegetables fermentation, salting, pickling.
5. With long storage, the content of nitrates in vegetables decreases.
But drying, juicing and mashing, on the contrary, increase the amount of nitrates.
1) cooking vegetables;
2) peeling;
3) removal of areas of the greatest accumulation of nitrates;
4) soaking.
In order to assess how real the danger of nitrate poisoning is, students are offered a calculation task.
A task. Table beet contains an average of 1200 mg of nitrate ions per 1 kg. When cleaning beets, 10% of nitrates are lost, and during cooking, another 40%. Will the daily intake of nitrates (325 mg) be exceeded if 200 g of boiled beets are eaten daily?
Given:
m(beets) = 1 kg,
With(NO 3 -) \u003d 1200 mg / kg,
m max (NO 3 - per day) = 325 mg,
m(beets) \u003d 200 g (0.2 kg),
(loss during cleaning) = 10%,
(loss during cooking) = 40%.
__________________________________
Find: m(NO 3 - in 200 g of boiled beets).
Solution
1 kg of beets - 1200 mg NO 3 -,
0.2 kg beets - X mg NO 3 -.
From here X= 240 mg (NO 3 -).
Total loss of nitrate ions:
(LOSS NO 3 -) = 10% + 40% = 50%.
Consequently, half of 240 mg or 120 mg of NO 3 - enters the body.
Answer. After cleaning and boiling the beets, the daily norm for nitrates (325 mg) contained in 200 g of the finished product (120 mg NO 3 -) is not exceeded, it can be eaten.
Nitrates in the production of explosives
Teacher. Many explosive mixtures contain an oxidizing agent (metal or ammonium nitrates, etc.) and fuel (diesel fuel, aluminum, wood flour). Therefore, salts - potassium nitrate, barium nitrate, strontium nitrate and others - are used in pyrotechnics..
What nitrogen fertilizer, together with aluminum and charcoal, is part of the explosive mixture - ammonal?
Student. Ammonal also contains ammonium nitrate. The main reaction that occurs during the explosion:
3NH 4 NO 3 + 2Al 3N 2 + 6H 2 O + Al 2 O 3 + Q.
The high heat of combustion of aluminum increases the energy of the explosion. The use of ammonium nitrate in the composition of ammonal is based on its ability to decompose upon detonation with the formation gaseous substances:
2NH 4 NO 3 (cr.) \u003d 2N 2 + 4H 2 O + O 2.
In the hands of terrorists, explosives bring only suffering to peaceful people.
For six centuries, the dominance of black powder in military affairs continued. Now it is used as an explosive in mining, in pyrotechnics (rockets, fireworks), and also as hunting gunpowder. Black or black powder is a mixture of 75% potassium nitrate, 15% charcoal and 10% sulfur.
An experience. Burning black or black powder
Prepare black powder by mixing 7.5 g of potassium nitrate, 1 g of sulfur and 1.5 g of charcoal. Before mixing, each substance is ground in a porcelain mortar. During the demonstration of the experiment, the mixture is placed in a pile on an iron sheet and set on fire with a burning torch. The mixture burns, forming a cloud of smoke (thrust).
Teacher. What role does saltpeter play?
Student. Saltpeter acts as an oxidizing agent when heated:
The use of nitrates and nitrites in medicine
5th student. Silver nitrate AgNO 3 , which blackens fabric, paper, desks and hands (lapis), is used as an antimicrobial agent for the treatment of skin ulcers, for cauterization of warts(the teacher demonstrates the technique of cauterizing warts on his hand) and as an anti-inflammatory agent for chronic gastritis and stomach ulcers: patients are prescribed to drink a 0.05% solution AgNO 3 . Powdered metals Zn, Mg, Al, mixed with silver nitrate, used in firecrackers.
Basic bismuth nitrate Bi (OH) 2 NO 3 is prescribed orally for peptic ulcer of the stomach and duodenum as an astringent and antiseptic agent. Outwardly - in ointments, powders for inflammatory skin diseases.
Salt sodium nitrite NaNO 2 used in medicine as an antispasmodic.
The use of nitrites in the food industry
6th student. Nitrites are used in sausage production: 7 g per 100 kg of minced meat. Nitrites give the sausage a pink color, without them it is gray, like boiled meat, and does not have a marketable appearance. In addition, the presence of nitrites in the sausage is necessary for another reason: they prevent the development of microorganisms that produce toxic poisons..
Knowledge control using the test "Nitrogen and its compounds"
Option I
1. The strongest molecule
a) H 2; b) F 2 ; c) O 2; d) N 2 .
2. Coloring of phenolphthalein in ammonia solution:
a) raspberry; b) green;
c) yellow; d) blue.
3. The oxidation state is +3 at the nitrogen atom in the compound:
a) NH 4 NO 3; b) NaNO 3 ; c) NO 2; d) KNO 2.
4. During thermal decomposition of copper (II) nitrate, the following are formed:
a) copper (II) nitrite and O 2;
b) nitric oxide (IV) and O 2 ;
c) copper oxide (II), brown gas NO 2 and O 2;
d) copper (II) hydroxide, N 2 and O 2.
5. Which ion is formed by the donor-acceptor mechanism?
a) ; b) NO 3 - ; c) Cl - ; d) SO 4 2–.
6. Specify strong electrolytes:
a) nitric acid;
b) nitrous acid;
c) an aqueous solution of ammonia;
d) ammonium nitrate.
7. Hydrogen is released during the interaction:
a) Zn + HNO 3 (razb.);
b) Cu + HCl (solution);
c) Al + NaOH + H 2 O;
d) Zn + H 2 SO 4 (razb.);
e) Fe + HNO 3 (conc.).
8. Write an equation for the reaction of zinc with very dilute nitric acid if one of the reaction products is ammonium nitrate. Specify the coefficient in front of the oxidizing agent.
9.
Name the substances A, B, C.
Option II
1. The water displacement method cannot be collected:
a) nitrogen; b) hydrogen;
c) oxygen; d) ammonia.
2. The reagent for the ammonium ion is a solution:
a) potassium sulfate; b) silver nitrate;
c) sodium hydroxide; d) barium chloride.
3. When HNO 3 (conc.) interacts with copper shavings, a gas is formed:
a) N 2 O; b) NH 3; c) NO 2 ; d) H 2 .
4. Thermal decomposition of sodium nitrate produces:
a) sodium oxide, brown gas NO 2, O 2;
b) sodium nitrite and O 2;
c) sodium, brown gas NO 2, O 2;
d) sodium hydroxide, N 2, O 2.
5. The degree of nitrogen oxidation in ammonium sulfate:
a) -3; b) -1; c) +1; d) +3.
6. With which of the following substances does concentrated HNO 3 react under normal conditions?
a) NaOH; b) AgCl; c) Al; d) Fe; e) Cu.
7. Specify the number of ions in the abbreviated ionic equation for the interaction of sodium sulfate and silver nitrate:
a) 1; b) 2; at 3; d) 4.
8. Write an equation for the interaction of magnesium with dilute nitric acid if one of the reaction products is a simple substance. Specify the coefficient in the equation in front of the oxidizing agent.
9. Write reaction equations for the following transformations:
Name the substances A, B, C, D.
Answers to test questions
Option I
1 - G; 2 - a; 3 - G; 4 - in; 5 - a; 6 - a, d; 7 - c, d; 8 – 10,
9. A - NH 3, B - NH 4 NO 3, C - NO,
Option II
1 - G; 2 - in; 3 - in; 4 – b; 5 - a; 6 - a, e; 7 - in,
2Ag + + SO 4 2– = Ag 2 SO 4;
8 – 12,
9. A - NO, B - NO 2, C - HNO 3, D - NH 4 NO 3,
At the end of the lesson, the teacher expresses his attitude to the work done by the students, evaluates their performances and answers.
LITERATURE
Gabrielyan O.S.. Chemistry-9. M.: Bustard, 2001; Gabrielyan O.S., Ostroumov I.G.. Handbook of the teacher. Chemistry. Grade 9 Moscow: Bustard, 2002; Pichugina G.V.. Generalization of knowledge about the transformation of nitrogen compounds in the soil and in plants. Chemistry at school, 1997, No. 7; Kharkivskaya N.L.,
Lyashenko L.F., Baranova N.V.. Beware of nitrates! Chemistry at school, 1999, No. 1; Zheleznyakova Yu.V., Nazarenko V.M.. Educational and research environmental projects. Chemistry at school, 2000, No. 3.
*“Red precipitate” is one of the modifications of mercury(II) oxide HgO. ( Note. ed.)
Nitrous acid is a monobasic weak acid that can only exist in dilute blue aqueous solutions and in gaseous form. Salts of this acid are called nitrites or nitrites. They are toxic and more stable than the acid itself. The chemical formula of this substance looks like this: HNO2.
Physical properties:
1. Molar mass equal to 47 g/mol.
2. is equal to 27 a.m.u.
3. Density is 1.6.
4. The melting point is 42 degrees.
5. The boiling point is 158 degrees.
Chemical properties of nitrous acid
1. If a solution with nitrous acid is heated, the following chemical reaction will occur:
3HNO2 (nitrous acid) \u003d HNO3 (nitric acid) + 2NO is released as a gas) + H2O (water)
2. Dissociates in aqueous solutions and is easily displaced from salts by stronger acids:
H2SO4 ( sulphuric acid) + 2NaNO2 (sodium nitrite) = Na2SO4 (sodium sulfate) + 2HNO2 (nitrous acid)
3. The substance we are considering can exhibit both oxidizing and reducing properties. When exposed to stronger oxidizing agents (for example: chlorine, hydrogen peroxide H2O2, oxidizes to nitric acid (in some cases, a salt of nitric acid is formed):
Restorative properties:
HNO2 (nitrous acid) + H2O2 (hydrogen peroxide) = HNO3 (nitric acid) + H2O (water)
HNO2 + Cl2 (chlorine) + H2O (water) = HNO3 (nitric acid) + 2HCl (hydrochloric acid)
5HNO2 (nitrous acid) + 2HMnO4 \u003d 2Mn (NO3) 2 (manganese nitrate, nitric acid salt) + HNO3 (nitric acid) + 3H2O (water)
Oxidizing properties:
2HNO2 (nitrous acid) + 2HI = 2NO (oxygen oxide, as gas) + I2 (iodine) + 2H2O (water)
Obtaining nitrous acid
This substance can be obtained in several ways:
1. When dissolving nitrogen oxide (III) in water:
N2O3 (nitric oxide) + H2O (water) = 2HNO3 (nitrous acid)
2. When dissolving nitrogen oxide (IV) in water:
2NO3 (nitric oxide) + H2O (water) = HNO3 (nitric acid) + HNO2 (nitrous acid)
Application of nitrous acid:
- diazotization of aromatic primary amines;
- production of diazonium salts;
- in the synthesis of organic substances (for example, for the production of organic dyes).
The effect of nitrous acid on the body
This substance is toxic, has a bright mutagenic effect, since in essence it is a deaminating agent.
What are nitrites
Nitrites are various salts of nitrous acid. They are less resistant to temperature than nitrates. Needed in the production of some dyes. Used in medicine.
Sodium nitrite has gained particular importance for humans. This substance has the formula NaNO2. Used as a preservative in Food Industry in the production of fish and meat products. It is a powder of pure white or slightly yellowish color. Sodium nitrite is hygroscopic (with the exception of purified sodium nitrite) and highly soluble in H2O (water). In air, it is able to gradually oxidize to have strong reducing properties.
Sodium nitrite is used in:
- chemical synthesis: to obtain diazo-amine compounds, to deactivate excess sodium azide, to obtain oxygen, sodium oxide and sodium nitrogen, to absorb carbon dioxide;
- in production food products(food additive E250): as an antioxidant and antibacterial agent;
- in construction: as an antifreeze additive to concrete in the manufacture of structures and building products, in the synthesis of organic substances, as an inhibitor of atmospheric corrosion, in the production of rubbers, poppers, additive solution for explosives; when processing metal to remove the tin layer and during phosphating;
- in photography: as an antioxidant and reagent;
- in biology and medicine: vasodilator, antispasmodic, laxative, bronchodilator; as an antidote for animal or human poisoning with cyanide.
Other salts of nitrous acid (eg potassium nitrite) are also currently used.
Nitrous acid
HNO 2 is a weak monobasic acid that exists only in dilute aqueous solutions.
Salts of nitrous acid are called nitrites. Nitrites are much more stable than HNO 2 and are all toxic.
Receipt:
1. N 2 O 3 + H 2 O \u003d 2HNO 2
How else can you get nitrous acid? ()
What is the oxidation state in nitrous acid?
This means that the acid exhibits both oxidizing and reducing properties.
Under the action of stronger oxidizing agents, it is oxidized to HNO 3:
5HNO 2 + 2HMnO 4 → 2Mn(NO 3) 2 + HNO 3 + 3H 2 O;
HNO 2 + Cl 2 + H 2 O → HNO 3 + 2HCl.
2HNO 2 + 2HI → 2NO + I 2 ↓ + 2H 2 O - reducing properties
Qualitative reaction to nitrite ion NO 2 – – interaction of nitrites with a solution of potassium iodide KI acidified with dilute sulfuric acid.
How should starch iodine paper change color under the action of free I 2?
Getting salts (nitrates and nitrites)
What are the methods of obtaining salts that you know? How can you get nitrates and nitrites?
1) Metal + non-metal = salt;
2) metal + acid = salt + hydrogen;
3) metal oxide + acid = salt + water;
4) metal hydroxide + acid = salt + water;
5) metal hydroxide + acid oxide = salt + water;
6) metal oxide + non-metal oxide = salt;
7) salt 1 + metal hydroxide (alkali) = salt 2 + metal hydroxide (insoluble base);
8) salt 1 + acid (strong) = salt 2 + acid (weak);
9) salt 1 + salt 2 = salt 3 + salt 4
10) salt 1 + metal (active) = salt 2 + metal (less active).
A specific way to obtain nitrates and nitrites:
disproportionation.
In the presence of excess oxygen
Salts of nitric acid - nitrates
nitrates of alkali metals, calcium, ammonium - saltpeter
KNO 3 - potassium nitrate,
NH 4 NO 3 - ammonium nitrate.
Physical properties:
All nitrates are solid crystalline substances, white in color, highly soluble in water. Poisonous!
Chemical properties of nitrates
Interaction of nitrates with metals, acids, alkalis, salts
Exercise. Mark the signs of each reaction, write down the molecular and ionic equations corresponding to the schemes:
Cu(NO 3) 2 + Zn…,
AgNO 3 + HCl ...,
Cu(NO 3) 2 + NaOH…,
AgNO 3 + BaCl 2 ....
Decomposition of nitrates
When solid nitrates are heated, they all decompose with the release of oxygen (the exception is ammonium nitrate), while they can be divided into three groups.
The first group consists of alkali metal nitrates
2KNO 3 \u003d 2KNO 2 + O 2.
The second group from alkaline earth metals up to and including copper
2Cu (NO 3) 2 \u003d 2CuO + 4NO 2 + O 2,
The third Me group after Cu
Hg (NO 3) 2 \u003d Hg + 2NO 2 + O 2,
Why is there a lot of nitrogen in nature (it is part of the atmosphere), and plants often give a poor harvest due to nitrogen starvation? (Plants cannot absorb molecular nitrogen from the air. With a lack of nitrogen, the formation of chlorophyll is delayed, the growth and development of the plant is delayed.)
Name the ways of assimilation of atmospheric nitrogen.
(Part of the bound nitrogen enters the soil during thunderstorms. Legumes, on the roots of which nodule bacteria develop that can bind atmospheric nitrogen, converting it into compounds available to plants.)
When harvesting, a person annually carries away with them huge amounts of bound nitrogen. He covers this loss by introducing not only organic, but also mineral fertilizers (nitrate, ammonia, ammonium). Nitrogen fertilizers are applied to all crops. Nitrogen is absorbed by plants in the form of ammonium cation and nitrate anion NO 3 - .
Student reports
Effect of nitrates on environment and the human body
First aid for nitrate poisoning
Reasons for the accumulation of nitrates in vegetables and methods for growing environmentally friendly crop products
HNO3, an oxygen-containing monobasic strong acid. Solid nitric acid forms two crystalline modifications with monoclinic and rhombic lattices.
Nitric acid is miscible with water in any ratio. In aqueous solutions, it almost completely dissociates into ions.
Obtained by the catalytic oxidation of synthetic ammonia on platinum-rhodium catalysts (Haber method) to a mixture of nitrogen oxides (nitrous gases), with their further absorption by water
4NH3 + 5O2 (Pt) > 4NO + 6H2O
2NO + O2 > 2NO2 4NO2 + O2 + 2H2O > 4HNO3 The concentration of nitric acid obtained by this method varies, depending on the technological design of the process, from 45 to 58%. For the first time, nitric acid was obtained by alchemists by heating a mixture of saltpeter and iron sulfate:
4KNO3 + 2(FeSO4 7H2O) (t°) > Fe2O3 + 2K2SO4 + 2HNO3^ + NO2^ + 13H2O
Pure nitric acid was first obtained by Johann Rudolf Glauber, acting on saltpeter with concentrated sulfuric acid:
KNO3 + H2SO4(conc.) (t°) > KHSO4 + HNO3^
Further distillation can be obtained so-called. "fuming nitric acid", containing practically no water.
Application:
in the production of mineral fertilizers;
in the military industry;
in photography - acidification of some tinting solutions;
in easel graphics - for etching printing forms (etching boards, zincographic printing forms and magnesium clichés).
1. Diluted nitric acid exhibits all the properties of strong acids; in aqueous solutions, it dissociates according to the following scheme:
HNO3 H+ + NO3–,
anhydrous acid:
2HNO3® NO2+ + NO3–+ H2O.
Gradually, especially in the light or when heated, nitric acid decomposes; during storage, the solution becomes brownish due to nitrogen dioxide:
4HNO3 4NO2 + 2H2O + O2.
2. Nitric acid interacts with almost all metals. Diluted nitric acid with alkali and alkaline earth metals, as well as with iron and zinc, forms the corresponding nitrates, ammonium nitrate or nitrogen hemioxide, depending on the activity of the metal, and water:
4Mg + 10HNO3® 4Mg(NO3)2 + N2O + 5H2O,
With heavy metals, dilute acid forms the corresponding nitrates, water, and nitric oxide is released, and in the case of a stronger dilution, nitrogen:
5Fe + 12HNO3(very diluted)®5Fe(NO3)3 + N2+ 6H2O,
3Cu + 8HNO3® 3Cu(NO3)2 + 2NO + 4H2O.
Concentrated nitric acid, when interacting with alkali and alkali metals, forms the corresponding nitrates, water, and nitrogen hemioxide is released:
8Na + 10HNO3® 8NaNO3 + N2O + 5H2O.
Concentrated acid passivates such metals as iron, chromium, aluminum, gold, platinum, iridium, tantalum, i.e. an acid-impermeable oxide film forms on the metal surface. Other heavy metals when interacting with concentrated nitric acid, they form the corresponding nitrates, water, and nitrogen oxide or dioxide is released:
3Hg + 8HNO3(cold)®3Hg(NO3)2 + 2NO + 4H2O,
Hg + 4HNO3(gor.)®Hg(NO3)2 + 2NO2+ 2H2O,
Ag + 2HNO3® AgNO3 + NO2+ 2H2O.
3. Nitric acid is able to dissolve gold, platinum and other precious metals, but mixed with hydrochloric acid. Their mixture in relation to three volumes of concentrated hydrochloric acid and one volume of concentrated nitric acid is called "aqua regia". The action of aqua regia is that nitric acid oxidizes hydrochloric acid to free chlorine, which combines with metals:
HNO3 + HCl ® Cl2 + 2H2O + NOCl,
2NOCl ® 2NO + Cl2.
Royal vodka is able to dissolve gold, platinum, rhodium, iridium and tantalum, which do not dissolve in nitrogen, and even more so hydrochloric acid:
Au + HNO3 + 3HCl ® AuCl3 + NO + 2H2O,
HCl + AuCl3® H;
3Pt + 4HNO3 + 12HCl ® 3PtCl4 + 4NO + 8H2O,
2HCl + PtCl4® H2.
4. Non-metals are also oxidized by nitric acid to the corresponding acids, dilute acid releases nitric oxide:
3P + 5HNO3 + 2H2O ® 3H3PO4 + 5NO ,
concentrated acid releases nitrogen dioxide:
S + 6HNO3® H2SO4 + 6NO2+ 2H2O,
nitric acid can also oxidize some inorganic compounds:
3H2S + 8HNO3® 3H2SO4 + 8NO + 4H2O.
HNO2 is a weak monobasic acid that exists only in dilute aqueous solutions, colored in a faint blue color, and in the gas phase. Salts of nitrous acid are called nitrites or nitrites. Nitrates are much more stable than HNO2, they are all toxic.
In the gas phase, the planar nitrous acid molecule exists in two configurations, cis- and trans-. At room temperature, the trans isomer predominates.
Chem. saints
In aqueous solutions, there is an equilibrium:
2HNO2 - N2O3 + H2O - NO^ + NO2^ + H2O
When the solution is heated, nitrous acid decomposes with the release of NO and NO2:
3HNO2 - HNO3 + 2NO^ + H2O.
HNO2 is a little stronger acetic acid. Easily displaced by stronger acids from salts:
H2SO4 + Ba(NO2)2 > BaSO4v + HNO2.
Nitrous acid exhibits both oxidizing and reducing properties. Under the action of stronger oxidizing agents (H2O2, KMnO4) it is oxidized to HNO3:
2HNO2 + 2HI > 2NO^ + I2v + 2H2O;
5HNO2 + 2HMnO4 >2Mn(NO3)2 + HNO3 + 3H2O;
HNO2 + Cl2 + H2O > HNO3 + 2HCl.
Nitrous acid is used to diazotize primary aromatic amines and form diazonium salts. Nitrites are used in organic synthesis in the production of organic dyes.
Receipt:
N2O3 + H2O 2HNO2,
NaNO2 + H2SO4 (0° C)® NaHSO4 + HNO2
AgNO2 + HCl ® AgCl + HNO2
Salt properties
All nitrates are highly soluble in water. With increasing temperature, their solubility increases greatly. When heated, nitrates decompose with the release of oxygen. Nitrates of ammonium, alkali and alkaline earth metals are called saltpeters, for example NaNO3 - sodium nitrate (Chilean nitrate), KNO3 - potassium nitrate, NH4NO3 - ammonium nitrate. Nitrates are obtained by the action of nitric acid HNO3 on metals, oxides, hydroxides, salts. Almost all nitrates are highly soluble in water.
Nitrates are stable at ordinary temperatures. They usually melt at relatively low temperatures (200-600°C), often with decomposition.
Alkali metal nitrates decompose to nitrites with the release of oxygen (and upon prolonged heating, they decompose stepwise into metal oxide, molecular nitrogen and oxygen, which is why they are good oxidizing agents).
Metal nitrates of medium activity decompose when heated to metal oxides with the release of nitrogen dioxide and oxygen.
Nitrates of the most inactive metals (noble metals) decompose mainly to free metals with the release of nitrogen dioxide and oxygen.
Nitrates are fairly strong oxidizing agents in solid state(usually in the form of a melt), but practically do not have oxidizing properties in solution, unlike nitric acid.
Nitrite is a salt of nitrous acid HNO2. Nitrites are thermally less stable than nitrates. They are used in the production of azo dyes and in medicine.
Nitric acid. Pure nitric acid HNO 3 is a colorless liquid with a density of 1.51 g / cm at - 42 ° C, solidifying into a transparent crystalline mass. In the air, it, like concentrated hydrochloric acid, "smokes", since its vapors form small droplets of fog with "moisture in the air,
Nitric acid does not differ in strength, Already under the influence of light, it gradually decomposes:
The higher the temperature and the more concentrated acid, the faster the decomposition. The released nitrogen dioxide dissolves in the acid and gives it a brown color.
Nitric acid is one of the strongest acids; in dilute solutions, it completely decomposes into H + and - NO 3 ions.
Oxidizing properties of nitric acid. A characteristic property of nitric acid is its pronounced oxidizing ability. Nitric acid-one
of the most energetic oxidizers. Many non-metals are easily oxidized by it, turning into the corresponding acids. So, when sulfur is boiled with nitric acid, it gradually oxidizes into sulfuric acid, phosphorus into phosphoric acid. A smoldering ember immersed in concentrated HNO 3 flares up brightly.
Nitric acid acts on almost all metals (with the exception of gold, platinum, tantalum, rhodium, iridium), turning them into nitrates, and some metals into oxides.
Concentrated HNO 3 passivates some metals. Lomonosov also discovered that iron, which dissolves easily in dilute nitric acid, does not dissolve.
in cold concentrated HNO 3 . Later it was found that nitric acid has a similar effect on chromium and aluminum. These metals go under
the action of concentrated nitric acid in a passive state.
The degree of oxidation of nitrogen in nitric acid is 4-5. Acting as an oxidizing agent, HNO 3 can be reduced to various products:
Receipt.
1. In the laboratory, nitric acid is obtained by reacting anhydrous nitrates with concentrated sulfuric acid:
Ba (NO 3) 2 + H 2 SO 4 → BaSO 4 ↓ + 2HNO 3.
2. In industry, the production of nitric acid goes in three stages:
1. Oxidation of ammonia to nitric oxide (II):
4NH 3 + 5O 2 → 4NO + 6 H 2 O
2. Oxidation of nitric oxide (II) to nitric oxide (IV):
2NO + O 2 → 2NO 2
3. Dissolution of nitric oxide (IV) in water with excess oxygen:
4NO 2 + 2H 2 O + O 2 → 4HNO 3
Chemical properties . Shows all the properties of acids. Nitric acid is one of the strongest mineral acids.
1. In aqueous solutions, it is completely dissociated into ions:
HNO 3 → H + + NO - 3
2. Reacts with metal oxides:
MgO + 2HNO 3 → Mg (NO 3) 2 + H 2 O,
3. Reacts with bases:
Mg (OH) 2 + 2HNO 3 → Mg (NO 3) 2 + 2H 2 O,
4. Concentrated HNO 3, when interacting with the most active metals to Al, is reduced to N 2 O. For example:
4Ca + 10HNO 3 → 4Ca(NO 3) 2 + N 2 O+ 5H 2 O
5. Concentrated HNO 3 when interacting with less active metals (Ni, Cu, Ag, Hg) is reduced to NO 2. For example:
4HNO 3 + Ni → Ni(NO 3) 2 + 2NO 2 + 2H 2 O.
6. Similarly, concentrated HNO 3 reacts with non-metals. The non-metal is oxidized. For example:
5HNO 3 + Po → HP + 5O 3 + 5NO 2 + 2H 2 O.
C nitric acid olis - nitrates when heated, they decompose according to the scheme:
to the left of Mg: MeNO 3 → MeNO 2 + O 2
Mg - Cu: MeNO 3 → MeO + NO 2 + O 2
to the right Cu MeNO 3 → Me + NO 2 + O 2
Application.
Nitric acid is used to produce nitrogen fertilizers, medicinal and explosives.
Hydrogen. The structure of the atom, physical and chemical properties, the production and use of hydrogen.
HYDROGEN, H, chemical element with atomic number 1, atomic mass 1,00794.
Natural hydrogen consists of a mixture of two stable nuclides with mass numbers 1.007825 (99.985% in the mixture) and 2.0140 (0.015%). In addition, in natural hydrogen there are always negligible amounts of a radioactive nuclide - tritium 3 H (half-life T1 / 2 = 12.43 years). Since the nucleus of a hydrogen atom contains only 1 proton (there cannot be less protons in the nucleus of an atom), it is sometimes said that hydrogen forms the natural lower boundary of the periodic system of elements of D. I. Mendeleev (although the element hydrogen itself is located in the uppermost part tables). The element hydrogen is located in the first period of the periodic table. It belongs to both the 1st group (group IA of alkali metals) and the 7th group (group VIIA of halogens).
The masses of atoms in hydrogen isotopes differ greatly (by several times). This leads to noticeable differences in their behavior in physical processes (distillation, electrolysis, etc.) and to certain chemical differences (differences in the behavior of isotopes of one element are called isotope effects; for hydrogen, isotope effects are most significant). Therefore, unlike the isotopes of all other elements, hydrogen isotopes have special symbols and names. Hydrogen with a mass number of 1 is called light hydrogen, or protium (lat. Protium, from the Greek protos - the first), denoted by the symbol H, and its nucleus is called a proton, symbol p. Hydrogen with a mass number of 2 is called heavy hydrogen, deuterium (Latin Deuterium, from the Greek deuteros - the second), the symbols 2 H, or D (read "de") are used to designate it, the nucleus d is the deuteron. A radioactive isotope with a mass number of 3 is called superheavy hydrogen, or tritium (lat. Tritum, from the Greek tritos - the third), the symbol 3 H or T (read "those"), the nucleus t is a triton.
The configuration of the only electron layer of the neutral unexcited hydrogen atom is 1s1. In compounds, it exhibits oxidation states +1 and, less often, -1 (valency I). The radius of the neutral hydrogen atom is 0.0529 nm. The ionization energy of the atom is 13.595 eV, the electron affinity is 0.75 eV. On the Pauling scale, the electronegativity of hydrogen is 2.20. Hydrogen is one of the non-metals.
In its free form, it is a light, flammable gas without color, odor or taste.
Physical and Chemical properties: under normal conditions, hydrogen is a light (density under normal conditions 0.0899 kg / m 3) colorless gas. Melting point -259.15°C, boiling point -252.7°C. Liquid hydrogen (at the boiling point) has a density of 70.8 kg/m 3 and is the lightest liquid. The standard electrode potential H 2 / H– in an aqueous solution is taken equal to 0. Hydrogen is poorly soluble in water: at 0 ° C, the solubility is less than 0.02 cm 3 / ml, but it is highly soluble in some metals (sponge iron and others), especially good - in metallic palladium (about 850 volumes of hydrogen in 1 volume of metal). The heat of combustion of hydrogen is 143.06 MJ/kg.
Exists in the form of diatomic H 2 molecules. The dissociation constant of H2 into atoms at 300 K is 2.56 10–34. The dissociation energy of the H 2 molecule into atoms is 436 kJ/mol. The internuclear distance in the H 2 molecule is 0.07414 nm.
Since the nucleus of each H atom that is part of the molecule has its own spin, molecular hydrogen can be in two forms: in the form of orthohydrogen (o-H 2) (both spins have the same orientation) and in the form of parahydrogen (p-H 2 ) (backs have different orientations). Under normal conditions, normal hydrogen is a mixture of 75% o-H 2 and 25% p-H 2 . The physical properties of p- and o-H 2 differ slightly from each other. Thus, if the boiling point pure o-n 2 20.45 K, then pure p-n 2 - 20.26 K. Turning o-n 2 in p-H 2 is accompanied by the release of 1418 J/mol of heat.
The high strength of the chemical bond between atoms in the H 2 molecule (which, for example, using the molecular orbital method, can be explained by the fact that in this molecule the electron pair is in the bonding orbital, and the loosening orbital is not populated with electrons) leads to the fact that at room temperature gaseous hydrogen is chemically inactive. So, without heating, with simple mixing, hydrogen reacts (with an explosion) only with gaseous fluorine (F):
H 2 + F 2 \u003d 2HF + Q.
If a mixture of hydrogen and chlorine (Cl) at room temperature is irradiated with ultraviolet light, then an immediate formation of hydrogen chloride HCl is observed. The reaction of hydrogen with oxygen (O) occurs with an explosion if a catalyst is added to the mixture of these gases - metallic palladium (Pd) (or platinum (Pt)). When ignited, a mixture of hydrogen and oxygen (O) (so-called explosive gas) explodes, and an explosion can occur in mixtures in which the hydrogen content is from 5 to 95 volume percent. Pure hydrogen in air or in pure oxygen (O) burns quietly with the release of a large amount of heat:
H 2 + 1 / 2O 2 \u003d H 2 O + 285.75 kJ / mol
If hydrogen interacts with other non-metals and metals, then only under certain conditions (heating, high pressure, the presence of a catalyst). So, hydrogen reacts reversibly with nitrogen (N) at elevated pressure (20-30 MPa and more) and at a temperature of 300-400 ° C in the presence of a catalyst - iron (Fe):
3H 2 + N 2 = 2NH 3 + Q.
Also, only when heated, hydrogen reacts with sulfur (S) to form hydrogen sulfide H 2 S, with bromine (Br) - to form hydrogen bromide HBr, with iodine (I) - to form hydrogen iodide HI. Hydrogen reacts with coal (graphite) to form a mixture of hydrocarbons of various compositions. Hydrogen does not interact directly with boron (B), silicon (Si), phosphorus (P), compounds of these elements with hydrogen are obtained indirectly.
When heated, hydrogen is able to react with alkali, alkaline earth metals and magnesium (Mg) to form compounds with an ionic bond nature, which contain hydrogen in the oxidation state –1. So, when calcium is heated in a hydrogen atmosphere, a salt-like hydride of the composition CaH 2 is formed. Polymeric aluminum hydride (AlH 3) x - one of the strongest reducing agents - is obtained indirectly (for example, using organoaluminum compounds). With many transition metals (for example, zirconium (Zr), hafnium (Hf), etc.), hydrogen forms compounds of variable composition (solid solutions).
Hydrogen is able to react not only with many simple, but also with complex substances. First of all, it should be noted the ability of hydrogen to reduce many metals from their oxides (such as iron (Fe), nickel (Ni), lead (Pb), tungsten (W), copper (Cu), etc.). So, when heated to a temperature of 400-450 ° C and above, iron (Fe) is reduced by hydrogen from any of its oxides, for example:
Fe 2 O 3 + 3H 2 \u003d 2Fe + 3H 2 O.
It should be noted that only metals located in the series can be reduced by hydrogen from oxides. standard potentials behind manganese (Mn). More active metals(including manganese (Mn)) are not reduced to metal from oxides.
Hydrogen is capable of adding to a double or triple bond to many organic compounds (these are the so-called hydrogenation reactions). For example, in the presence of a nickel catalyst, hydrogenation of ethylene C 2 H 4 can be carried out, and ethane C 2 H 6 is formed:
C 2 H 4 + H 2 \u003d C 2 H 6.
The interaction of carbon monoxide (II) and hydrogen in industry produces methanol:
2H 2 + CO \u003d CH 3 OH.
In compounds in which a hydrogen atom is connected to an atom of a more electronegative element E (E \u003d F, Cl, O, N), hydrogen bonds form between the molecules (two E atoms of the same or two different elements are interconnected through the H atom: E "... N ... E"", with all three atoms located on the same straight line). Such bonds exist between the molecules of water, ammonia, methanol, etc. and lead to a noticeable increase in the boiling points of these substances, an increase in the heat of evaporation and etc.
Receipt: Hydrogen can be obtained in many ways. In industry, natural gases are used for this, as well as gases obtained from oil refining, coking and gasification of coal and other fuels. In the production of hydrogen from natural gas (the main component is methane), its catalytic interaction with water vapor and incomplete oxidation with oxygen (O) are carried out:
CH 4 + H 2 O \u003d CO + 3H 2 and CH 4 + 1/2 O 2 \u003d CO 2 + 2H 2
The separation of hydrogen from coke gas and refinery gases is based on their liquefaction during deep cooling and removal from the mixture of gases that are more easily liquefied than hydrogen. In the presence of cheap electricity, hydrogen is obtained by electrolysis of water, passing current through alkali solutions. Under laboratory conditions, hydrogen is easily obtained by the interaction of metals with acids, for example, zinc (Zn) with hydrochloric acid.
Application: hydrogen is used in the synthesis of ammonia NH3, hydrogen chloride HCl, methanol CH 3 OH, in the hydrocracking (cracking in a hydrogen atmosphere) of natural hydrocarbons, as a reducing agent in the production of certain metals. By hydrogenation of natural vegetable oils, hard fat is obtained - margarine. Liquid hydrogen finds use as a rocket fuel and also as a coolant. A mixture of oxygen (O) and hydrogen is used in welding.
At one time, it was suggested that in the near future, the reaction of hydrogen combustion will become the main source of energy production, and hydrogen energy will replace traditional sources of energy production (coal, oil, etc.). At the same time, it was assumed that for the production of hydrogen on a large scale it would be possible to use the electrolysis of water. Water electrolysis is a rather energy-intensive process, and it is currently unprofitable to obtain hydrogen by electrolysis on an industrial scale. But it was expected that electrolysis would be based on the use of medium temperature (500-600°C) heat, which occurs in large quantities during the operation of nuclear power plants. This heat is of limited use, and the possibility of obtaining hydrogen with its help would solve both the problem of ecology (when hydrogen is burned in air, the amount of environmentally harmful substances formed is minimal) and the problem of utilization of medium-temperature heat. However, after the Chernobyl disaster, the development of nuclear energy is curtailed everywhere, so that the indicated source of energy becomes inaccessible. Therefore, the prospects for the widespread use of hydrogen as an energy source are still shifting, at least until the middle of the 21st century.
Features of circulation : hydrogen is not poisonous, but when handling it, one must constantly take into account its high fire and explosion hazard, and the explosion hazard of hydrogen is increased due to the high ability of the gas to diffuse even through some solid materials. Before starting any heating operations in an atmosphere of hydrogen, you should make sure that it is clean (when igniting hydrogen in a test tube turned upside down, the sound should be dull, not barking).
27 The position of microorganisms in the system of the living world. Diversity of microorganisms and their commonality with other organisms. The essential features of microorganisms are: small cell size, high metabolic activity, high plasticity of their metabolism (rapid adaptation to changing environmental conditions, "ubiquity"), the ability to reproduce rapidly, poor morphological differentiation, and a variety of metabolic processes.
Microorganisms, (microbes) - the collective name for a group of living organisms that are too small to be visible to the naked eye (their characteristic size is less than 0.1 mm). Microorganisms include both non-nuclear (prokaryotes: bacteria, archaea) and eukaryotes: some fungi, protists, but not viruses, which are usually isolated into a separate group. Most microorganisms consist of a single cell, but there are also multicellular microorganisms, just as there are some unicellular macroorganisms visible to the naked eye, such as Thiomargarita namibiensis, representatives of the genus Caulerpa (they are giant polykaryons). Microbiology is the study of these organisms.
The ubiquity and total power of the metabolic potential of microorganisms determines their most important role in the circulation of substances and maintaining dynamic balance in the Earth's biosphere.
A brief review of various representatives of the microcosm, occupying certain "floors" of size, shows that, as a rule, the size of objects is definitely related to their structural complexity. The lower size limit for a free-living single-celled organism is determined by the space required to pack inside the cell the apparatus necessary for independent existence. The limitation of the upper limit of the size of microorganisms is determined, according to modern concepts, by the relationship between the cell surface and volume. With an increase in cellular dimensions, the surface increases in the square, and the volume in the cube, so the ratio between these values shifts towards the latter.
Microorganisms live almost everywhere where there is water, including hot springs, the bottom of the world's oceans, and also deep inside the earth's crust. They are an important link in the metabolism in ecosystems, mainly acting as decomposers, but in some ecosystems they are the only producers of biomass.
Microorganisms that live in various environments, participate in the cycle of sulfur, iron, phosphorus and other elements, decompose organic substances of animal, vegetable origin, as well as abiogenic origin (methane, paraffins), provide self-purification of water in reservoirs.
However, not all types of microorganisms are beneficial to humans. A very large number of species of microorganisms is opportunistic or pathogenic for humans and animals. Some microorganisms cause damage to agricultural products, deplete the soil with nitrogen, cause pollution of water bodies, and the accumulation of toxic substances (for example, microbial toxins) in food products.
Microorganisms are characterized by good adaptability to the action of environmental factors. Various microorganisms can grow at temperatures from −6° to +50-75°. The record for survival at elevated temperature was set by archaea, some of the studied cultures of which grow on nutrient media above 110 ° C, for example, Methanopyrus kandleri (strain 116) grows at 122 ° C, a record high temperature for all known organisms.
In nature, habitats with this temperature exist under pressure in hot volcanic springs at the bottom of the oceans (Black smokers).
Microorganisms are known that thrive at levels of ionizing radiation that are fatal for multicellular creatures, in a wide range of pH values, at 25% sodium chloride concentration, in conditions of various oxygen contents up to its complete absence (Anaerobic microorganisms).
At the same time, pathogenic microorganisms cause diseases in humans, animals and plants.
The most widely accepted theories about the origin of life on Earth suggest that protomicroorganisms were the first living organisms to emerge through evolution.
Currently, all microorganisms are divided into 3 kingdoms:
1. Procariota. All types of bacteria, rickettsia, chlamydia, mycoplasmas, etc. can be attributed to this kingdom. Cells have a nucleus with one chromosome. The nucleus is not separated from the cytoplasm of the cell. A simple dividing cycle by constriction. There are a number of unique organelles such as plasmids, mesosomes. There is no ability for photosynthesis.
2. Eucariotae. Representatives of this kingdom are fungi and protozoa. The cell contains a nucleus, delimited from the cytoplasm by a membrane, with several chromosomes. There are a number of organelles characteristic of higher animals: mitochondria, endoplasmic reticulum, Golgi apparatus. Some representatives of this kingdom have chloroplasts and are capable of photosynthesis. They have a complex life cycle.
3. Vira. Viruses belong to this kingdom. The hallmark of a virion is the presence of only one type of nucleic acid: RNA or DNA enclosed in a capsid. A virus may not have a common outer shell. The reproduction of the virus can occur only after embedding in another cell, where replication takes place.
