Sulfurous acid is an inorganic dibasic unstable acid of medium strength. An unstable compound, known only in aqueous solutions at a concentration of no more than six percent. When attempting to isolate pure sulfurous acid, it breaks down into sulfur oxide (SO2) and water (H2O). For example, when concentrated sulfuric acid (H2SO4) reacts with sodium sulfite (Na2SO3), sulfur oxide (SO2) is released instead of sulfurous acid. This is what the reaction looks like:
Na2SO3 (sodium sulfite) + H2SO4 ( sulfuric acid) = Na2SO4 (sodium sulfate) + SO2 (sulfur dioxide) + H2O (water)
Sulfurous acid solution
When storing it, it is necessary to exclude access to air. Otherwise, sulfurous acid, slowly absorbing oxygen (O2), will turn into sulfuric acid.
2H2SO3 (sulfuric acid) + O2 (oxygen) = 2H2SO4 (sulfuric acid)
Solutions of sulfurous acid have a rather specific odor (reminiscent of the odor remaining after lighting a match), the presence of which can be explained by the presence of sulfur oxide (SO2), which is not chemically bound with water.
Chemical properties sulfurous acid
1. H2SO3) can be used as a reducing agent or an oxidizing agent.
H2SO3 is a good reducing agent. With its help, it is possible to obtain hydrogen halides from free halogens. For example:
H2SO3 (sulfuric acid) + Cl2 (chlorine, gas) + H2O (water) = H2SO4 (sulfuric acid) + 2HCl (hydrochloric acid)
But when interacting with strong reducing agents, this acid will act as an oxidizing agent. An example is the reaction of sulfurous acid with hydrogen sulfide:
H2SO3 (sulfuric acid) + 2H2S (hydrogen sulfide) = 3S (sulfur) + 3H2O (water)
2. The chemical compound we are considering forms two - sulfites (medium) and hydrosulfites (acidic). These salts are reducing agents, just like (H2SO3) sulfurous acid. When they are oxidized, salts of sulfuric acid are formed. When sulfites of active metals are calcined, sulfates and sulfides are formed. This is a self-oxidation-self-healing reaction. For example:
4Na2SO3 (sodium sulfite) = Na2S + 3Na2SO4 (sodium sulfate)
Sodium and potassium sulfites (Na2SO3 and K2SO3) are used in dyeing fabrics in the textile industry, in bleaching metals, and in photography. Calcium hydrosulfite (Ca(HSO3)2), which exists only in solution, is used to process wood material into a special sulfite pulp. It is then used to make paper.
Application of sulfurous acid
Sulfurous acid is used:
For bleaching wool, silk, wood pulp, paper and other similar substances that cannot withstand bleaching with stronger oxidizing agents (for example, chlorine);
As a preservative and antiseptic, for example, to prevent the fermentation of grain when producing starch or to prevent the fermentation process in wine barrels;
To preserve food, for example, when canning vegetables and fruits;
Processed into sulfite pulp, from which paper is then produced. In this case, a solution of calcium hydrosulfite (Ca(HSO3)2) is used, which dissolves lignin, a special substance that binds cellulose fibers.
Sulfurous acid: preparation
This acid can be produced by dissolving sulfur dioxide (SO2) in water (H2O). You will need concentrated sulfuric acid (H2SO4), copper (Cu) and a test tube. Algorithm of actions:
1. Carefully pour concentrated sulfuric acid into a test tube and then place a piece of copper in it. Heat up. The following reaction occurs:
Cu (copper) + 2H2SO4 (sulfuric acid) = CuSO4 (sulfur sulfate) + SO2 (sulfur dioxide) + H2O (water)
2. The flow of sulfur dioxide must be directed into a test tube with water. When it dissolves, it partially occurs with water, resulting in the formation of sulfurous acid:
SO2 (sulfur dioxide) + H2O (water) = H2SO3
So, by passing sulfur dioxide through water, you can get sulfurous acid. It is worth considering that this gas has an irritating effect on the membranes of the respiratory tract, can cause inflammation, as well as loss of appetite. Inhaling it for a long time may cause loss of consciousness. This gas must be handled with extreme caution and care.
Slide 2
Sulfur
Sulfur is a chalcogen, a fairly reactive non-metal. There are three allotropic modifications of sulfur: orthorhombic S8 plastic monoclinic
Slide 3
Characteristics of sulfur
Serav PSHE: position (period, group) atomic structure properties of the element by period / in the main p/g higher oxide higher hydroxide LVS
Slide 4
Receipt
When draining solutions of hydrogen sulfide and sulfurous acids: H2SO3 + 2H2S = 3S + 3H2O When hydrogen sulfide is incompletely burned (with a lack of air): 2H2S + O2 = 2S + 2H2O
Slide 5
Chemical properties
Does not wet and does not react with water. The oxidizing agent reacts with: metals (except gold) Hg + S = HgS (neutralization of spilled mercury) hydrogen and non-metals, whose s.o. less (carbon, phosphorus, etc.)
Slide 6
How does a reducing agent react with: oxygen, chlorine, fluorine
Slide 7
S-2(with me, C, P, H2): C + 2S = CS2 H2 + S = H2S S0 S S+2 S + Cl2 = SCl2 S+4 S + O2 = SO2H2SO3 S+6 S + 3F2 = SF6H2SO4 strengthening the oxidizing ability of ions
Slide 8
Hydrogen sulfide
H2S – hydrogen sulfide. Its solution in water is called hydrosulfide acid. The acid is weak dibasic, therefore it dissociates stepwise: I: H2S ↔ H+ + HS– II: HS– ↔ H+ + S–
Slide 9
Shows all the properties of acids. Reacts with: basic oxides: H2S + CaO = CaS + H2O bases: H2S + KOH ↔ KHS + H2O H2S + OH– ↔ HS– + H2O H2S + 2KOH ↔ K2S + H2O H2S + 2OH– ↔ S2– + H2O
Slide 10
salts: CuCO3 + H2S = CuS + H2CO3 metals: Ca + H2S = CaS + H2
Slide 11
Properties of salts
Acid salts of hydrosulfide acid - hydrosulfides (KHS, NaHS) are highly soluble in water. Sulfides of alkali and alkaline earth metals are also soluble. Sulfides of other metals are insoluble in water, and sulfides of copper, lead, silver, mercury and other heavy metals are insoluble even in acids (except nitric).
Slide 12
Hydrogen sulfide oxidation
Hydrogen sulfide is easily oxidized by oxygen (as with an excess of O2 and a deficiency?). Bromine water Br2: H2S + Br2 = 2HBr + S↓ yellow-orange colorless
Slide 13
Sulfur(IV) oxide
SO2 – sulfur dioxide gas. Reacts with water to form H2SO3. Typical acid oxide. Reacts with bases (salt (sulfite or hydrosulfite) and water are formed) and basic oxides (only salt is formed).
Slide 14
Obtained by: burning sulfur, roasting pyrite, the action of acids on sulfites, the interaction of conc. sulfuric acid and heavy meth
Slide 15
Sulfur(VI) oxide
SO3 is an acidic oxide. It reacts with water to form H2SO4, with bases (salt (sulfate or hydrosulfate) and water is formed) and basic oxides. Obtained by oxidation of sulfur dioxide. Dissolves in sulfuric acid to form oleum: H2SO4 + nSO3 = H2SO4 nSO3 oleum
Slide 16
Sulfuric acid
Sulfuric acid H2SO4 is a heavy, odorless, colorless oily liquid. At a concentration > 70%, sulfuric acid is called concentrated, less than 70% - dilute. The dissociation of sulfuric acid is expressed by the equation: H2SO4 ↔ 2H++ SO42–
Slide 17
The acid reacts with amophoteric and basic oxides and hydroxides, salts: H2SO4 + BaCl2 = BaSO4↓ + HCl The latter reaction is qualitative for SO42–ion (an insoluble white precipitate is formed).
Slide 18
H2SO4 H2SO4 +1 +6 -2 H2SO4 +1 +6 -2 diluted concentrated H+ ― oxidizing agent 2H+ + 2e– = H2 S+6 ― oxidizing agent S+6 +8e– +6e– +2e– S-2 (H2S) S0 (S)S+4 (SO2)
Slide 19
All metals in the activity series up to hydrogen react with dilute sulfuric acid. During the reaction, metal sulfate is formed and hydrogen is released: H2SO4 + Zn = ZnSO4 + H2 Metals after hydrogen do not react with dilute acid: Cu + H2SO4 ≠
Slide 20
Concentrated sulfuric acid
Metals in the activity series after hydrogen interact with concentrated sulfuric acid according to the following scheme: H2SO4 (conc.) + Me = MeSO4 + SO2 + H2O I.e. formed: metal sulfate sulfur(IV) oxide - sulfur dioxide SO2 water
Slide 21
More active mesulfuric acid under certain conditions can be reduced to sulfur in pure form or hydrogen sulfide. In the cold conc. sulfuric acid passivates iron and aluminum, so they are transported in iron tanks: H2SO4 (conc.) + Fe ≠ (in the cold)
Slide 22
Preparation of sulfuric acid
obtaining SO2 (usually by roasting pyrite) oxidation of SO2 into SO3 in the presence of a catalyst - vanadium(V) oxide; dissolution of SO3 in sulfuric acid to obtain oleum
Slide 23
Sulfates
Salts of sulfuric acid have all the properties of salts. Their relationship to heating is special: active metal sulfates (Na, K, Ba) do not decompose even at t > 1000˚C others (Cu, Al, Fe) even with slight heating decompose into sulfur oxide (VI) and metal oxide
Slide 24
Questions
In what reactions does sulfur play the role of an oxidizing agent? reducing agent? what degrees does it exhibit? What causes the difference in the properties of concentrated and dilute sulfuric acid? write the equations for the reaction of conc. and dilute acids with copper and zinc. how to distinguish solutions of sodium iodide and sodium sulfate? propose two methods and write the reaction equations in molecular and ionic forms.
Slide 25
Quests
How much sulfur dioxide can be obtained from 10 kg of ore containing 48% pyrite? What volume is occupied by: a) 4 moles of SO2? b) 128 g SO3? Carry out the reactions: O2 → S → SO2 → SO3 → H2SO4 → Na2SO4 → BaSO4
View all slides
O.S.ZAYTSEV
CHEMISTRY BOOK
FOR TEACHERS SECONDARY SCHOOLS,
STUDENTS OF PEDAGOGICAL UNIVERSITIES AND SCHOOLCHILDREN OF 9–10 GRADES,
WHO DECIDED TO DEVOTE THEMSELVES TO CHEMISTRY AND NATURAL SCIENCE
TEXTBOOK TASK LABORATORY PRACTICAL SCIENTIFIC STORIES FOR READING
Continuation. See No. 4–14, 16–28, 30–34, 37–44, 47, 48/2002;
1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 12, 13, 14, 15, 16, 17, 18, 19, 20, 21, 22, 23,
24, 25-26, 27-28, 29, 30, 31, 32, 35, 36, 37, 39, 41, 42, 43, 44 , 46, 47/2003;
1, 2, 3, 4, 5, 7, 11, 13, 14, 16, 17, 20, 22, 24/2004
§ 8.1. Redox reactions
LABORATORY RESEARCH
(continuation)
2. Ozone is an oxidizing agent.
Ozone is the most important substance for nature and humans.
Ozone creates an ozonosphere around the Earth at an altitude of 10 to 50 km with a maximum ozone content at an altitude of 20–25 km. Being in the upper layers of the atmosphere, ozone does not allow most of the ultraviolet rays of the Sun, which have a detrimental effect on humans, animals and flora. IN recent years Areas of the ozonosphere with greatly reduced ozone content, the so-called ozone holes, have been discovered. It is not known whether ozone holes have formed before. The reasons for their occurrence are also unclear. It is believed that chlorine-containing freons from refrigerators and perfume cans are ultraviolet radiation The sun releases chlorine atoms, which react with ozone and thereby reduce its concentration in the upper atmosphere. Scientists are extremely concerned about the danger of ozone holes in the atmosphere.
In the lower layers of the atmosphere, ozone is formed as a result of a series of sequential reactions between atmospheric oxygen and nitrogen oxides emitted by poorly adjusted car engines and discharges from high-voltage power lines. Ozone is very harmful to breathing - it destroys the tissue of the bronchi and lungs. Ozone is extremely toxic (more powerful than carbon monoxide). The maximum permissible concentration in the air is 10–5%.
Thus, ozone in the upper and lower layers of the atmosphere has opposite effects on humans and the animal world.
Ozone, along with chlorine, is used to treat water to break down organic impurities and kill bacteria. However, both chlorination and ozonation of water have their advantages and disadvantages. When water is chlorinated, bacteria are almost completely destroyed, but organic substances of a carcinogenic nature that are harmful to health (promote the development of cancer) are formed - dioxins and similar compounds. When water is ozonized, such substances are not formed, but ozone does not kill all bacteria, and after some time the remaining living bacteria multiply abundantly, absorbing the remains of killed bacteria, and the water becomes even more contaminated with bacterial flora. Therefore, ozonation of drinking water is best used when it is used quickly. Ozonation of water in swimming pools is very effective when water continuously circulates through the ozonizer. Ozone is also used for air purification. It is one of the environmentally friendly oxidizing agents that do not leave harmful products of its decomposition.
Ozone oxidizes almost all metals except gold and platinum group metals.
Chemical methods ozone production is ineffective or too dangerous. Therefore, we advise you to obtain ozone mixed with air in an ozonizer (the effect of a weak electrical discharge on oxygen) available in the school physics laboratory.
Ozone is most often obtained by acting on gaseous oxygen with a quiet electrical discharge (without glow or sparks), which occurs between the walls of the internal and external vessels of the ozonator. The simplest ozonizer can be easily made from glass tubes with stoppers. You will understand how to do this from Fig. 8.4. The inner electrode is a metal rod (long nail), the outer electrode is a wire spiral. Air can be blown out with an aquarium air pump or a rubber bulb from a spray bottle. In Fig. 8.4 The inner electrode is located in a glass tube ( Why do you think?), but you can assemble an ozonizer without it. Rubber plugs are quickly corroded by ozone.
 |
High voltage It is convenient to obtain from the induction coil of the car's ignition system by continuously opening the connection to the low voltage source (battery or 12 V rectifier).
The ozone yield is several percent.
Ozone can be detected qualitatively using a starch solution of potassium iodide. A strip of filter paper can be soaked in this solution, or the solution can be added to ozonized water, and air with ozone can be passed through the solution in a test tube. Oxygen does not react with iodide ion.
Reaction equation:
2I – + O 3 + H 2 O = I 2 + O 2 + 2OH – .
Write the equations for the reactions of electron gain and loss.
Bring a strip of filter paper moistened with this solution to the ozonizer. (Why does potassium iodide solution need to contain starch?) Hydrogen peroxide interferes with the determination of ozone using this method. (Why?).
Calculate the EMF of the reaction using the electrode potentials:
3. Reductive properties of hydrogen sulfide and sulfide ion.
Hydrogen sulfide is a colorless gas with the smell of rotten eggs (some proteins contain sulfur).
To conduct experiments with hydrogen sulfide, you can use gaseous hydrogen sulfide, passing it through a solution with the substance being studied, or add pre-prepared hydrogen sulfide water to the solutions under study (this is more convenient). Many reactions can be carried out with a solution of sodium sulfide (reactions with the sulfide ion S 2–).
Work with hydrogen sulfide only under draft! Mixtures of hydrogen sulfide with air burn explosively.
Hydrogen sulfide is usually produced in a Kipp apparatus by reacting 25% sulfuric acid (diluted 1:4) or 20% hydrochloric acid (diluted 1:1) on iron sulfide in the form of pieces 1–2 cm in size. Reaction equation:
FeS (cr.) + 2H + = Fe 2+ + H 2 S (g.).
Small quantities of hydrogen sulfide can be obtained by placing crystalline sodium sulfide in a stoppered flask through which a dropping funnel with a stopcock and an outlet tube are passed. Slowly pouring 5–10% from the funnel hydrochloric acid (why not sulfur?), the flask is constantly shaken by shaking to avoid local accumulation of unreacted acid. If this is not done, unexpected mixing of components can lead to a violent reaction, expulsion of the stopper and destruction of the flask.
A uniform flow of hydrogen sulfide is obtained by heating hydrogen-rich organic compounds, such as paraffin, with sulfur (1 part paraffin to 1 part sulfur, 300 ° C).
To obtain hydrogen sulfide water, hydrogen sulfide is passed through distilled (or boiled) water. About three volumes of hydrogen sulfide gas dissolve in one volume of water. When standing in air, hydrogen sulfide water gradually becomes cloudy. (Why?).
Hydrogen sulfide is a strong reducing agent: it reduces halogens to hydrogen halides, and sulfuric acid to sulfur dioxide and sulfur.
Hydrogen sulfide is poisonous. The maximum permissible concentration in the air is 0.01 mg/l. Even at low concentrations, hydrogen sulfide irritates the eyes and respiratory tract and causes headaches. Concentrations above 0.5 mg/l are life-threatening. At higher concentrations it is affected nervous system. Inhaling hydrogen sulfide may cause cardiac and respiratory arrest. Sometimes hydrogen sulfide accumulates in caves and sewer wells, and a person trapped there instantly loses consciousness and dies.
At the same time, hydrogen sulfide baths have a healing effect on the human body.
3a. Reaction of hydrogen sulfide with hydrogen peroxide.
Study the effect of hydrogen peroxide solution on hydrogen sulfide water or sodium sulfide solution.
Based on the results of the experiments, compose reaction equations. Calculate the EMF of the reaction and draw a conclusion about the possibility of its passage.
3b. Reaction of hydrogen sulfide with sulfuric acid.
Pour concentrated sulfuric acid dropwise into a test tube with 2–3 ml of hydrogen sulfide water (or sodium sulfide solution). (carefully!) until turbidity appears. What is this substance? What other products might be produced in this reaction?
Write the reaction equations. Calculate the emf of the reaction using electrode potentials:
4. Sulfur dioxide and sulfite ion.
Sulfur dioxide, sulfur dioxide, is the most important atmospheric pollutant emitted by automobile engines when using poorly purified gasoline and by furnaces in which sulfur-containing coals, peat or fuel oil are burned. Every year, millions of tons of sulfur dioxide are released into the atmosphere due to the burning of coal and oil.
Sulfur dioxide occurs naturally in volcanic gases. Sulfur dioxide is oxidized by atmospheric oxygen into sulfur trioxide, which, absorbing water (vapor), turns into sulfuric acid. Falling acid rain destroys cement parts of buildings, architectural monuments sculptures carved from stone. Acid rain slows down the growth of plants and even leads to their death, and kills living organisms in water bodies. Such rains wash out phosphorus fertilizers, which are poorly soluble in water, from arable lands, which, when released into water bodies, lead to rapid proliferation of algae and rapid swamping of ponds and rivers.
Sulfur dioxide is a colorless gas with a pungent odor. Sulfur dioxide should be obtained and worked with under draft.
Sulfur dioxide can be obtained by placing 5–10 g of sodium sulfite in a flask closed with a stopper with an outlet tube and a dropping funnel. From a dropping funnel with 10 ml concentrated sulfuric acid (extreme caution!) pour it drop by drop onto the sodium sulfite crystals. Instead of crystalline sodium sulfite, you can use its saturated solution.
Sulfur dioxide can also be produced by the reaction between copper metal and sulfuric acid. In a round-bottomed flask equipped with a stopper with a gas outlet tube and a dropping funnel, place copper shavings or pieces of wire and pour a little sulfuric acid from the dropping funnel (about 6 ml of concentrated sulfuric acid is taken per 10 g of copper). To start the reaction, warm the flask slightly. After this, add the acid drop by drop. Write the equations for accepting and losing electrons and the total equation.
The properties of sulfur dioxide can be studied by passing the gas through a reagent solution, or in the form of an aqueous solution (sulfurous acid). The same results are obtained when using acidified solutions of sodium sulfites Na 2 SO 3 and potassium sulfites K 2 SO 3. Up to forty volumes of sulfur dioxide are dissolved in one volume of water (a ~6% solution is obtained).
Sulfur dioxide is toxic. With mild poisoning, a cough begins, a runny nose, tears appear, and dizziness begins. Increasing the dose leads to respiratory arrest.
4a. Interaction of sulfurous acid with hydrogen peroxide.
Predict the reaction products of sulfurous acid and hydrogen peroxide. Test your assumption with experience.
Add the same amount of 3% hydrogen peroxide solution to 2–3 ml of sulfurous acid. How to prove the formation of the expected reaction products?
Repeat the same experiment with acidified and alkaline solutions of sodium sulfite.
Write the reaction equations and calculate the emf of the process.
Select the electrode potentials you need:
4b. Reaction between sulfur dioxide and hydrogen sulfide.
This reaction takes place between gaseous SO 2 and H 2 S and serves to produce sulfur. The reaction is also interesting because the two air pollutants mutually destroy each other. Does this reaction take place between solutions of hydrogen sulfide and sulfur dioxide? Answer this question with experience.
Select electrode potentials to determine whether a reaction can occur in solution:
Try to carry out a thermodynamic calculation of the possibility of reactions. The thermodynamic characteristics of substances to determine the possibility of a reaction between gaseous substances are as follows:
In which state of substances - gaseous or in solution - are reactions more preferable?
Lesson 22 9th grade
Lesson on: Hydrogen sulfide. Sulfides. Sulfur oxide (IV). Sulfurous acid
Lesson objectives: General education: To consolidate students' knowledge on the topic covered: allotropy of sulfur and oxygen, the structure of sulfur and oxygen atoms, chemical properties and use of sulfur using testing, in order to prepare students for the State Examination; Study the structure, properties and use of gases: hydrogen sulfide, sulfur dioxide, sulfurous acid. Study salts - sulfides, sulfites and their qualitative determination using an electronic textbook in chemistry for grade 9. Study the influence of hydrogen sulfide, sulfur oxide (IV) on the environment and human health. Use student presentations when studying new topic and consolidation. Use a multimedia projector when checking the test. Continue preparing students for passing chemistry exams in the form of State Examinations.
Educational: Moral and aesthetic education students to the environment. Developing a belief in the positive role of chemistry in life modern society, the need for a chemically literate attitude towards your health and the environment. Developing the ability to work in pairs during self-analysis of control sections and tests.
Educational: Be able to apply acquired knowledge to explain a variety of chemical phenomena and properties of substances. Be able to apply additional material from information sources, computer technologies in preparing students for the State Examination. Use the acquired knowledge and skills in practical activities And everyday life: a) environmentally conscious behavior in the environment; b) impact assessments chemical pollution environment on the human body.
Equipment for the lesson: G.E. Rudzitis, F.G. Feldman "Chemistry textbook 9th grade." Student presentations: “Hydrogen sulfide”, “Sulfur oxide (IV)", "Ozone". Test for preparing the State Examination Test, answers to the test. Electronic manual for studying chemistry, grade 9: a) qualitative reactions to sulfide ion, sulfite ion. b) multimedia projector
c) projection screen. Protection of the poster “Environmental pollution by emissions of hydrogen sulfide and sulfur dioxide.”
Progress of the lesson.
I. Beginning of the lesson: The teacher announces the topic, purpose and objectives of the lesson.
Consolidation of the studied material:
It is carried out on test questions in order to prepare students for passing the State Examination Test (test attached).
The test answers are displayed on the screen:
Students mutually check the tests and give grades (the sheets are handed over to the teacher).Evaluation criteria: 0 errors – 5; 1 – 2 errors – 4; 3 errors – 3; 4 and more – 2
The test is carried out within 7 minutes and checked within 3 minutes.
II. Learning a new topic:
Hydrogen sulfide. Sulfides.
Hydrogen sulfide is a chemically valuable sulfur compound; we will study its properties in today’s lesson. We will get acquainted with the presence of hydrogen sulfide in nature, its physical properties and its effect on the human body and the environment through a presentation.
Why is it impossible to obtain hydrogen sulfide in the laboratory like other gases, for example: oxygen and hydrogen? Students will answer this question after listening to the presentation.
Structure of hydrogen sulfide:
a) molecular formula H 2 S -2 , oxidation state of sulfur (-2), toxic.
b) hydrogen sulfide has the smell of rotten eggs.
3. Preparation of hydrogen sulfide: Preparation in the laboratory: obtained by the action of dilute sulfuric acid on iron sulfide (II), since hydrogen sulfide is poisonous, experiments are carried out in a fume hood.H 2 + S 0 → H 2 S -2
FeS + H 2 SO 4 → FeSO 4 + H 2 Sthis reaction is carried out in a Kip apparatus, which is used to produce hydrogen.
4. Chemical properties of hydrogen sulfide: Hydrogen sulfide burns in air with a blue flame and produces sulfur dioxide or sulfur oxide (IV)
2 H 2 S -2 + 3 O 2 → 2 H 2 O + 2 S +4 O 2
reducing agent
With a lack of oxygen, water and sulfur vapors are formed: 2H 2 S -2 + O 2 → 2 H 2 O + 2 S 0
Hydrogen sulfide has the properties of a reducing agent: if a small amount of bromine water is added to a test tube with hydrogen sulfide, the solution will become discolored and sulfur will appear on the surface of the solution
H 2 S -2 + Br 0 2 → S 0 + 2 HBr -1
Hydrogen sulfide is slightly soluble in water: in one volume of water att= 20 º 2.4 volumes of hydrogen sulfide dissolve, this solution is called hydrogen sulfide water or weak hydrogen sulfide acid. Consider the dissociation of hydrosulfide acid:H 2 S→ H + +HS -
H.S. - ↔ H + + S 2- Dissociation in the second step practically does not occur, since it is a weak acid. It gives 2 types of salts:
H.S. - (I)S 2-
hydrosulfides sulfides
IIIII
NaHSNa 2 S
Sodium hydrosulfide sodium sulfide
Hydrogen sulfide acid reacts with alkalis in a neutralization reaction:
H 2 S + NaOH → NaHS + H 2 O
excess
H 2 S+2NaOH→ Na 2 S+2H 2 O
excess
Qualitative reaction to sulfide ion (demonstration of experience from an electronic educational disk)
Pb(NO 3 ) 2 + Na 2 S → PbS↓ + 2 NaNO 3 write a complete ionic and a short one
black precipitate ionic equation
(Na 2 S + CuCl 2 → CuS↓ + 2 HCl)
black sediment
Exercise for the eyes. (1-2 minutes)
Compliance with sanitary and hygienic standards for working with a computer in the classroom.
5. Sulfur oxide( IV) – sulfur dioxide.S +4 O 2 sulfur oxidation degree (+4).
Another important sulfur compound is sulfur oxide (IV) SO 2 – sulfur dioxide. Poisonous.
WITH physical properties We will get acquainted with sulfur dioxide, its application and impact on the environment and human health through a presentation.
Why sulfur dioxide cannot be obtained from practical work?
Obtaining sulfur oxide(IV): formed when sulfur burns in air, a gas with a pungent odor.
S+O 2 → SO 2
Sulfur dioxide has the properties of an acidic oxide; when dissolved in water, sulfurous acid is formed, an electrolyte of medium strength.SO 2 + H 2 O ↔ H 2 SO 3 litmus turns red.
Chemical propertiesSO 2 :
Reacts with basic oxidesSO 2 + CaO → CaSO 3
Reacts with alkalisSO 2 + 2 NaOH → Na 2 SO 3 + H 2 O
(at home, write down the complete ionic equation and the short ionic equation)
Sulfur exhibits oxidation states:S -2 , S 0 , S +4 , S +6 .
In sulfur oxide( IV) SO 2 oxidation state +4, therefore sulfur dioxide exhibits the properties of an oxidizing agent and a reducing agent
S +4 O 2 + 2H 2 S -2 → 3S 0 ↓ + 2H 2 O S +4 O 2 +Cl 0 2 + 2H 2 O → H 2 S +6 O 4 + 2HCl -1 2-
Hydrosulfite sulfite
TO HSO 3 K 2 SO 3
Qualitative reaction to sulfite ion (the reagent is sulfuric acid, a gas with a pungent odor is formed that discolors solutions) fragment from an electronic educational disk.
K 2 SO 3 + H 2 SO 4 → K 2 SO 4 + SO 2 + H 2 O
At home, write down the complete and short ionic equation.
Protection of the poster “Environmental pollution with sulfur compounds.”
Protecting your presentation
Homework §11-12, notes, ex. 3.5 p.34(p)
III. Lesson summary:
The teacher sums up the lesson
Gives grades for tests and presentations.
Thanks students for the lesson.
First aid for gas poisoning: hydrogen sulfide, sulfur dioxide: rinsing the nose and mouth with a 2% solution of sodium bicarbonateNaHCO 3 , peace, fresh air.
The physical properties of sulfur directly depend on the allotropic modification. For example, the most famous modification of sulfur is rhombic, S₈. It is quite fragile crystalline substance yellow color.
Structure of the rhombic sulfur molecule S₈
In addition to the rhombic one, there are many other modifications, the number of which, according to different sources, reaches three dozen.
Chemical properties of the element
At normal temperatures, the chemical activity of sulfur is quite low. But when heated, sulfur often interacts with all simple substances, metals and non-metals.
S + O₂ → SO₂
Sulfur is an essential element in life and animals, and is widely used in industries ranging from medicine to pyrotechnic devices.
Sulfuric acid
Sulfuric acid has the formula H₂SO₄ and is the strongest dibasic acid. Previously, this substance was called oil of vitriol because the concentrated acid has a thick, oily consistency.
Sulfuric acid mixes easily with water, but such solutions must be prepared with caution: concentrated acid must be carefully poured into water, and in no case vice versa.
Sulfuric acid is a caustic substance and can dissolve some. Therefore, it is often used in ore mining. Acid leaves severe burns on the skin, so it is extremely important to follow safety precautions when working with it.
Obtaining "oil of vitriol"
The industry uses a contact method for producing SO₂ (sulfur dioxide) through the oxidation of sulfur dioxide, which is formed during the combustion of sulfur. Next, sulfur trioxide SO₃ is obtained from sulfur dioxide, which is then dissolved in the most concentrated sulfuric acid. The resulting solution is called oleum. To obtain “oil of vitriol,” oleum is diluted with water.
Chemical properties of sulfuric acid
When interacting with metals, as well as carbon and sulfur, concentrated sulfuric acid oxidizes them:
Сu + 2H₂SO₄ (conc.) → CuSO₄ + SO₂ + 2H₂O.
C(graphite) + 2H₂SO₄ (conc., horizontal) → CO₂ + 2SO₂ + 2H₂O
S + 2H₂SO₄ (conc.) → 3SO₂ + 2H₂O
Dilute acid is capable of reacting with all metals that are to the left of hydrogen in the voltage series:
Fe + H₂SO₄ (dil.) → FeSO₄ + H₂
Zn + H₂SO₄ (dil.) → ZnSO₄ + H₂
In reactions with bases, dilute H₂SO₄ forms sulfates and hydrosulfates:
H₂SO₄ + NaOH → NaHSO₄ + H₂O;
H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.
This acid can also react with basic oxides, resulting in sulfates:
CaO + H₂SO₄ → CaSO₄↓ + H₂O.