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 vegetable world. IN last 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.
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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. Reducing 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% hydrochloric acid from the funnel (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, 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?
Hydrogen sulfide (H₂S) is a colorless gas with a rotten egg odor. It is denser than hydrogen. Hydrogen sulfide is deadly poisonous to humans and animals. Even a small amount of it in the air causes dizziness and nausea, but the worst thing is that after inhaling it for a long time, this smell is no longer felt. However, for hydrogen sulfide poisoning, there is a simple antidote: you should wrap a piece of bleach in a handkerchief, then moisten it, and sniff the package for a while. Hydrogen sulfide is produced by reacting sulfur with hydrogen at a temperature of 350 °C:
H₂ + S → H₂S
This is a redox reaction: during it, the oxidation states of the elements participating in it change.
In laboratory conditions, hydrogen sulfide is produced by treating iron sulfide with sulfuric or hydrochloric acid:
FeS + 2HCl → FeCl₂ + H₂S
This is an exchange reaction: in it, the interacting substances exchange their ions. This process is usually performed using a Kipp apparatus.
Kipp apparatus
Properties of hydrogen sulfide
When hydrogen sulfide burns, sulfur oxide 4 and water vapor are formed:
2H₂S + 3О₂ → 2Н₂О + 2SO₂
H₂S burns with a bluish flame, and if you hold an inverted beaker over it, clear condensate (water) will appear on its walls.
However, with a slight decrease in temperature, this reaction proceeds somewhat differently: a yellowish coating of free sulfur will appear on the walls of the pre-cooled glass:
2H₂S + O₂ → 2H₂O + 2S
The industrial method for producing sulfur is based on this reaction.
When a pre-prepared gaseous mixture of hydrogen sulfide and oxygen is ignited, an explosion occurs.
The reaction of hydrogen sulfide and sulfur(IV) oxide also produces free sulfur:
2H₂S + SO₂ → 2H₂O + 3S
Hydrogen sulfide is soluble in water, and three volumes of this gas can dissolve in one volume of water, forming weak and unstable hydrosulfide acid (H₂S). This acid is also called hydrogen sulfide water. As you can see, the formulas of hydrogen sulfide gas and hydrogen sulfide acid are written the same way.
If a solution of lead salt is added to hydrosulfide acid, a black precipitate of lead sulfide will form:
H₂S + Pb(NO₃)₂ → PbS + 2HNO₃
This is a qualitative reaction for the detection of hydrogen sulfide. It also demonstrates the ability of hydrosulfide acid to enter into exchange reactions with salt solutions. Thus, any soluble lead salt is a reagent for hydrogen sulfide. Some other metal sulfides also have a characteristic color, for example: zinc sulfide ZnS - white, cadmium sulfide CdS - yellow, copper sulfide CuS - black, antimony sulfide Sb₂S₃ - red.
By the way, hydrogen sulfide is an unstable gas and, when heated, almost completely decomposes into hydrogen and free sulfur:
H₂S → H₂ + S
Hydrogen sulfide interacts intensively with aqueous solutions of halogens:
H₂S + 4Cl₂ + 4H₂O→ H₂SO₄ + 8HCl
Hydrogen sulfide in nature and human activity
Hydrogen sulfide is part of volcanic gases, natural gas and gases associated with oil fields. There is a lot of it in natural mineral waters, for example, in the Black Sea it lies at a depth of 150 meters and below.
Hydrogen sulfide is used:
- in medicine (treatment with hydrogen sulfide baths and mineral waters);
- in industry (production of sulfur, sulfuric acid and sulfides);
- in analytical chemistry (for the precipitation of heavy metal sulfides, which are usually insoluble);
- in organic synthesis (to produce sulfur analogues of organic alcohols (mercaptans) and thiophene (sulphur-containing aromatic hydrocarbon). Another recently emerging area in science is hydrogen sulfide energy. The production of energy from hydrogen sulfide deposits from the bottom of the Black Sea is being seriously studied.
The nature of redox reactions of sulfur and hydrogen
The reaction of hydrogen sulfide formation is redox:
Н₂⁰ + S⁰→ H₂⁺S²⁻
The process of interaction of sulfur with hydrogen is easily explained by the structure of their atoms. Hydrogen occupies first place in the periodic table, therefore, the charge of its atomic nucleus is equal to (+1), and 1 electron circles around the atomic nucleus. Hydrogen easily gives up its electron to atoms of other elements, turning into a positively charged hydrogen ion - a proton:
Н⁰ -1е⁻= Н⁺
Sulfur is in position sixteen in the periodic table. This means that the charge of the nucleus of its atom is (+16), and the number of electrons in each atom is also 16e⁻. The location of sulfur in the third period suggests that its sixteen electrons swirl around the atomic nucleus, forming 3 layers, the last of which contains 6 valence electrons. The number of valence electrons of sulfur corresponds to the number of group VI in which it is located in the periodic table.
So, sulfur can give up all six valence electrons, as in the case of the formation of sulfur(VI) oxide:
2S⁰ + 3O2⁰ → 2S⁺⁶O₃⁻²
In addition, as a result of the oxidation of sulfur, 4e⁻ can be given up by its atom to another element to form sulfur(IV) oxide:
S⁰ + O2⁰ → S⁺4 O2⁻²
Sulfur can also donate two electrons to form sulfur(II) chloride:
S⁰ + Cl2⁰ → S⁺² Cl2⁻
In all three of the above reactions, sulfur donates electrons. Consequently, it is oxidized, but at the same time acts as a reducing agent for oxygen atoms O and chlorine Cl. However, in the case of the formation of H2S, oxidation is the lot of hydrogen atoms, since they are the ones who lose electrons, restoring the external energy level sulfur from six electrons to eight. As a result, each hydrogen atom in its molecule becomes a proton:
Н2⁰-2е⁻ → 2Н⁺,
and the sulfur molecule, on the contrary, being reduced, turns into a negatively charged anion (S⁻²): S⁰ + 2е⁻ → S⁻²
Thus, in chemical reaction In the formation of hydrogen sulfide, the oxidizing agent is sulfur.
From the point of view of the manifestation of sulfur in various oxidation states, another interesting interaction between sulfur(IV) oxide and hydrogen sulfide is the reaction to produce free sulfur:
2H₂⁺S-²+ S⁺⁴О₂-²→ 2H₂⁺O-²+ 3S⁰
As can be seen from the reaction equation, both the oxidizing agent and the reducing agent in it are sulfur ions. Two sulfur anions (2-) donate two of their electrons to the sulfur atom in the sulfur(II) oxide molecule, as a result of which all three sulfur atoms are reduced to free sulfur.
2S-² - 4е⁻→ 2S⁰ - reducing agent, oxidizes;
S⁺⁴ + 4е⁻→ S⁰ - oxidizing agent, reduced.
Lesson 13
Sulfur(IV) oxide. Hydrogen sulfide and sulfurous acid and their salts
Lesson objectives:
1. Characterize the chemical properties of sulfur oxide (IV), hydrogen sulfide and sulfurous acids and their salts, qualitative reactions to sulfur compounds(subject result).
2. Continue to develop the ability to generate ideas, identify cause-and-effect relationships, look for analogies and work in a team, use alternative sources of information(metasubject result).
3. Formation of skills to manage your educational activities, preparation for understanding the choice of further educational trajectory(personal result).
During the classes
Preparing to perceive new material (10 min)
Survey of students on homework.
Learning new material (20 min)
Hydrogen sulfide H 2 S – colorless gas, heavier than air, smells like rotten eggs. Very poisonous. Contained in volcanic gases and mineral waters.
Obtained by exchange reaction:
1. Burning in air with a blue flame:
2H 2 S+3O 2( hut .) = 2H 2 O+2SO 2
2H 2 S+O 2( lack .) = 2H 2 O+2S
2. Restorative properties:
3. When dissolved in water, hydrosulfide acid is formed, which dissociates:
4. Interaction with alkalis. Forms two types of salts: sulfides and hydrosulfides:
Sulfur dioxide SO 2 : colorless, with a pungent odor, heavier than air, soluble in water, poisonous.
Acidic oxide.
1. When mixed with water, it forms a sulfurous compound:
Sulfurous acid unstable, easily decomposes into sulfur oxide (IV) and water. Exists only in aqueous solutions. Forms two types of salts: sulfites and hydrosulfites.
Qualitative reaction to sulfites
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 of 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.