The use of research tasks in chemistry lessons. Research work of students at school in chemistry Obtaining ozone with an electric discharge in oxygen
Research at home in the kitchen under the guidance of a teacher Research objectives: Educational: to provide additional information about acids and bases, to use them correctly; formation of report writing skills; to teach students to think independently, find and solve problems. Developing: develop the ability to highlight the main thing, generalize, classify; acquire knowledge independently. Educational: to teach to independently evaluate, observe phenomena; develop an interest in the subject and Creative skills in the process of independent work; the formation of interest in a new subject.
The report on the research work is being carried out according to plan. 1. The title of the topic of work. The title should accurately reflect the content of the work. Date, venue, last name and first name of the author. 2. The purpose of the work and its tasks. 3. Method of work. The results of the work depend on the number of experiments, observations and their processing. In what ways were observations made, how many of them were carried out, with what substances. 4. Results and discussion. Multiple students can receive the same assignment. Therefore, it is necessary to discuss the results of experiments, observations, comparison of reports.
Research methodology. one. Preparatory stage: For experiments, a small amount of vegetables, fruits, baking soda, vinegar, juices will be required, therefore, it is necessary to ask parents not to be sorry if the child spoils them in their experiments, because the child will learn the world, and this is a step into a great science. 2. Acquaintance with the object of study. The student receives a card - a task. 3. Familiarization with safety precautions.
Conducting research. Work 1. Acids and bases in the kitchen. You will need: vinegar, lemon juice, orange juice, apple juice, citric acid, sparkling water, baking soda, detergent, glasses. Pour a full spoonful of baking soda into an empty glass. Pour some vinegar into a glass. What do you observe?. Try lemon, orange, apple juice, soda, detergent. Mix drop detergent with any liquid acid (vinegar, fruit juice or soda). Add a small amount of the resulting mixture to a spoon with baking soda. Does this create foam? Foam formation indicates that the solution continues to be an acid. Add more detergent to the previously prepared mixture. Continue testing the acidic properties of the mixture by observing the foam release. Stopping the formation of foam will mean the neutralization of the acid.
Work 2. Growing crystals. You will need: salt, sugar, water, transparent plastic cups, spoon, rope, pencil. Place a few full tablespoons of table salt in a glass. Fill the glass three-quarters full with water. Stir the salt with a spoon. If the salt has dissolved, add another tablespoon of salt, stir and add salt until the solution is saturated. Tie the rope to the middle of the pencil, and lower the free end of the rope with a spoon to the bottom of the glass. The next day, you will see that crystals stand out on the walls of the glass and on the rope. Repeat the experiment using sugar or another salt. Leave the experimental setup for a week, thus allowing time for maximum crystallization to occur. Carefully study the formed crystals, and you will notice that they are of different shapes. Replace rope with thread. Separate a single crystal and observe it. Every day it will increase in size.
Work 3. Shiny coin. You will need: any copper-containing coin, salt, vinegar, paper towel, spoon. Place the coin on a paper towel. Sprinkle some salt on it. Pour the vinegar over the top with a spoon. Rub a coin and it will shine before your eyes! Repeat this experiment with a) one salt. b) one vinegar. c) with lemon juice. d) with salt and lemon juice. Does one of the above combinations clean the coin as effectively as using vinegar and salt?
Research lessons are becoming popular among chemistry teachers. Such lessons require a lot of preparation, which, as practice shows, justifies itself. Such lessons are built in accordance with the logic of the activity approach and include the following stages: motivational-orienting, operational-executive (analysis, forecasting and experiment), evaluative-reflective.
Conducting a thought experiment. Helps develop reasoning skills. These are tasks in which you need to get a specific substance from those offered; get the substance in several ways; carry out all the characteristic and qualitative reactions inherent in this class of substances; reveal the genetic relationship between the classes of inorganic substances.
Examples of tasks for a thought experiment. 1. Zinc powder was poured into the retort, the gas outlet tube was closed with a clamp, the retort was weighed and the contents were calcined. When the retort had cooled, it was weighed again. Has the mass changed and why? Then the clamp was opened. Has the mass changed and why? 2. Cups with solutions of sodium hydroxide and sodium chloride are balanced on the scales. Will the position of the arrow of the scales change after some time and why?
Creative tasks for predicting the properties of substances. Such tasks contribute to the formation of research skills, stimulate interest, allow students to get acquainted with the achievements of scientists, see beautiful, elegant vivid examples of the work of creative thought.
For example, when studying the topic Carbohydrates, students are asked questions: 1. German chemist Christian Schönbein accidentally spilled a mixture of sulfuric and nitric acids on the floor. He automatically wiped the floor with his wife's cotton apron. Acid can set fire to the apron, thought Shenbein, rinsed the apron in water and hung it up to dry over the stove. The apron dried up, but then there was a quiet explosion and ... the apron disappeared. Why did the explosion happen? 2. What happens if you chew bread crumb for a long time?
Lesson topic: Chemical properties of nitric acid. The general didactic goal of the lesson: to create conditions for the primary awareness and comprehension of educational information in order to develop the research skills of students by means of technology problem learning. Triune didactic goal: Educational aspect: to promote the formation of the concept of "acid" in students using the example of nitric acid; create conditions for identifying the general and specific properties of nitric acid by solving experimental and cognitive problems, develop skills in writing reaction equations. Developing aspect: to promote the development of students' research skills in the process of performing and observing the experiment. Educational aspect: to maintain interest in the study of the topic through independent work; foster cooperation; to promote the development of competent chemical speech.
Goals for students: be able to write reaction equations involving nitric acid in various situations and transfer the knowledge gained to solve practical problems; working at a creative level: to be able to analyze the conditions of processes, find various options for their solution, predict the results of the interaction of nitric acid with other substances. Type of lesson: learning new material. Teaching methods: partially exploratory, research, reproductive.
Forms of Methods Implementation: Problem Seminar. Techniques for implementing methods: creating tasks of a research nature; tasks for comparison and analysis of previously received information; assignments for independent transfer of knowledge to a new learning situation. Forms of organization of cognitive activity: general class, group (in this lesson it provides for facilitating the implementation of experimental research work, contributes to the creation of adaptive educational environment and economy of reagents), individual. Expected result: all students will learn the general and specific properties of nitric acid, as well as why a solution of nitric acid interacts with metals differently from solutions of other acids.
Pedagogical conclusions 1. Students of different levels of preparedness and different ages are included in research activities with pleasure and interest, i. the assertion that this is an area of interests and opportunities for high school students and that this type of activity is only possible for gifted children is incorrect. Teachers involved in research activities of students of different levels of preparedness should take into account the child's capabilities, predict the level of the result, the rate of implementation of the research program. 2. In the course of research activities, the development of the child's abilities occurs under certain conditions: - if the topic and subject of research activity correspond to the needs of the child; - training takes place in the zone of proximal development and for sufficient high level difficulties; - if the content of the activity is based on the subjective experience of the child; - if there is a learning of ways of activity. 3. Teaching research skills begins with a lesson that is built according to the laws of conducting scientific research. The technology of research activity is focused on the development of skills: - to determine the goals and objectives of the study, its subject; - independent search for literature and its notes; - analysis and systematization of information; - annotate the studied sources; - put forward a hypothesis, conduct a practical study in accordance with it from the classification of the material; - describe the results of the study, draw conclusions and generalizations.
One of the most important tasks of the teacher is the development of students' mental abilities (which is no less important than the simple acquisition of knowledge and skills), which is possible only in the process of independent creative search for new knowledge and ways of activity, i.e., in solving problems that arise in the course of research , I organized by the teacher.
As experience shows, teaching and research activities contribute to:
expanding and updating students' knowledge of the subjects of the school curriculum, developing their interest in the disciplines studied, as well as ideas about interdisciplinary connections;
development of the intellectual initiative of schoolchildren in the process of mastering basic and additional educational programs;
creation of prerequisites for the development of a scientific way of thinking among students;
mastering their creative approach to any kind of activity;
learning to use in their work Information Technology and other means of communication;
formation in the educational institution of a developing educational environment for the child;
professional self-determination of students;
receiving pre-professional training;
meaningful organization of children's free time.
When carrying out research activities on the basis of an experiment, the following stages of general scientific activity are assumed:
setting the goal of the experiment, which determines what result the experimenter intends to obtain in the course of the study;
formulation and substantiation of a hypothesis that can be used as the basis for an experiment. A hypothesis is a set of theoretical propositions, the truth of which is subject to verification;
experiment planning, which takes place in the following sequence: 1) drawing up a plan for the experiment and, if necessary, an image of the device design; thinking through the work after the end of the experiment (disposal of reagents, features of washing dishes, etc.); 2) selection of laboratory equipment and reagents; 3) identification of the source of danger (description of precautionary measures during the experiment); 4) the choice of the form of registration of the results of the experiment;
implementation of the experiment, fixing observations and measurements;
analysis, processing and explanation of the results of the experiment, which include: 1) mathematical processing of the results of the experiment (if necessary); 2) comparison of the results of the experiment with the hypothesis; 3) explanation of the ongoing processes in the experiment; 4) formulation of conclusions;
reflection - awareness and evaluation of the experiment based on a comparison of the goal and results, during which it is necessary to find out whether all the operations to carry out the experiment were successful.
Tasks constitute a special group heuristic and research character. In doing so, students use reasoning as a means of obtaining subjectively new knowledge about substances and chemical reactions. At the same time, schoolchildren carry out theoretical research, on the basis of which they formulate definitions, find relationships between the structure and properties, the genetic relationship of substances, systematize facts and establish patterns, conduct an experiment in order to solve the problem formed by the teacher or set independently.
For example, when studying the properties of amphoteric hydroxides, the following task can be proposed: “Will the result of the interaction of solutions of sodium hydroxide and aluminum chloride be the same when the first is added to the second and vice versa?”
When studying the topic “Generalization of the properties of the main classes of inorganic substances”, students can be asked to answer the question: “What happens if a solution of sodium hydroxide is added to a solution of copper (II) sulfate, and potassium hydroxide is added to a solution of sodium carbonate?”
On the topic "Halogens" questions may be of interest:
1. What color will an indicator paper acquire in a freshly prepared solution of chlorine in water?
2. What color will an indicator paper have in a chlorine solution that has been exposed to light for some time?
The answers to these questions are confirmed experimentally.
Practice shows that the use creative tasks, which consist in predicting the properties of substances, contributes to the formation of research skills, stimulates interest, allows students to get acquainted with the achievements of scientists, see beautiful, elegant vivid examples of the work of creative thought.
When studying the topic "Carbohydrates", students can complete the following tasks:
1. German chemist Christian Schönbein accidentally spilled a mixture of sulfuric and nitric acids on the floor. He automatically wiped the floor with his wife's cotton apron. "Acid can set fire to the apron," thought Schoenbein, rinsed the apron in water and hung it up to dry over the stove. The apron dried up, but then there was a quiet explosion and ... the apron disappeared. Why did the explosion happen?
2. What happens if you chew a piece of bread for a long time?
As a result of performing laboratory experiment No. 3 (grade 7) “Studying the signs of a chemical reaction (gas evolution),” students must make sure that the main sign in the interaction of chalk and acetic acid is gas evolution. However, more observant students may note another sign: the dissolution of the solid substance of chalk in acetic acid. To consolidate the results of the experience, students can be asked to answer the following questions:
1. Where at home do we meet with a similar sign of reaction?
2. What substance can be used instead of vinegar in the preparation of fizzy drinks?
In laboratory experiment No. 6 (grade 7) "Interaction of acids with metals", students receive experimental confirmation of a number of metal activity and a laboratory method for producing hydrogen. You can ask them to find answers to the following questions:
1. What other metals can be used to produce hydrogen from acids?
2. Why can't mercury be used to produce hydrogen?
When studying the topic “Proteins”, students can pose the following question: “Why can’t shoes made of genuine leather be dried on a central heating battery?”
To answer the question, students make a plan for finding the answer:
a) the protein composition of the skin;
b) the structure of the protein molecule;
c) the effect of temperature on the protein structure.
Then they find the answer: “High temperature, causing denaturation and destructuring of the protein, leads to a change in the strength and size of shoes.” A problematic demonstration experiment can also be used in the work, for example: testing substances and their solutions for electrical conductivity; reaction of ammonium salts with alkalis; neutralization of acids with ammonia (“smoke without fire”); interaction of metals with salt solutions; the ratio of aluminum to concentrated nitric acid; reaction of ethylene with bromine water and potassium permanganate solution; amphotericity of aluminum hydroxide; reaction of glycerol with copper(II) hydroxide, etc.
An experiment can be included either at the formulation stage or at the problem solving stage. In the latter case, the experience confirms (or does not confirm) the hypothesis put forward by the students, and the problem is determined using other methods and methods. In this case, a special role is played by thought experiment, which develops abstract thinking. These include tasks in which it is necessary to obtain a specific substance from the proposed ones; get it in several ways; mentally sort through all the characteristic and qualitative reactions inherent in this class of substances; reveal the genetic relationship between the classes of inorganic substances. It is impossible to neglect the thought experiment, it can be carried out at all stages of the lesson in the form of group, frontal or individual work.
For example, in a lesson on the topic “Halogens and their salts”, at the stage of fixing the material, instead of a reproductive question about the colors of silver halides, you can offer a thought experiment to recognize solutions of halides.
When studying the topic "Electrolytic dissociation", the traditional experimental determination of the electrical conductivity of substances using a device begins with a mental one! experiment. After that, it is carried out
demonstration experiment. Students compare and analyze the results, make drawings and diagrams in notebooks, write down the equations for the reaction of electrolytic dissociation. Examples of thought experiment tasks:
1. Zinc powder was poured into the retort, the gas outlet tube was closed with a clamp, the retort was weighed, and the contents were calcined. When the retort had cooled, it was weighed again. Has its mass changed and why?
2. Then the clamp was opened. Has the mass changed and why?
3. Cups with solutions of sodium hydroxide and sodium chloride are balanced on the scales. Will the position of the arrow of the scales change after some time and why?
4. Suggest methods for producing ethyl alcohol using natural gas and water as feedstock.
5. Make the equations for the reactions of obtaining acetic acid based on limestone, coal, water, air.
6. How can aniline be obtained if natural gas, air and water are used as feedstock?
7. Natural honey contains glucose and fructose. Suggest methods for obtaining artificial honey.
8. Suggest your ways of solving the problem of turning liquid fats into solid ones. What economically beneficial raw materials for Belarus can be used for this?
9. Select and justify the most cost-effective ways to obtain glycerin to soften the leather of winter boots.
10. Suggest methods for detecting in natural water or water that has passed an industrial water treatment system: a) excess acidity or alkalinity; b) ammonium cations; c) nitrate anions.
11. Do you have a suspicion that the employees of the gas station,
where your father constantly fills up the car, they add water to gasoline. You have quicklime at your disposal. Is it possible to check your suspicions with her help?
When studying qualitative reactions to ions, students acquire the ability to draw up a plan for the recognition of substances. The class is divided into groups of four people and each of them is given the task to draw up a plan for determining solutions of sulfate, carbonate and sodium chloride, which are in three numbered test tubes. Mandatory conditions: visibility. Desired conditions: speed and minimum spent reagents. Each group defends their plan, using previously acquired knowledge, writes down the molecular and ionic reaction equations. In conclusion, students conduct a laboratory experiment, putting their plan into practice.
A special place in the educational process is occupied by exercises that form in schoolchildren ideas about such a method of scientific research as modeling. The following tasks can help them with this:
1. Make models of atoms of oxygen, sulfur, selenium, tellurium. Compare their properties.
2. Based on the structure of atoms, determine the type of chemical bond in compounds H 2 S , H 2 O, H 2 Se . How does the polarization of a chemical bond change with an increase in the radius of the elements Group VI?
3.What is an ecological house? Suggest a model.
4.When cooking in the kitchen, there is a specific smell of acrolein aldehyde. Draw up a structural
the formula of this substance, if it is known that its molecular formula is C 3 H 4 0 and
aldehyde is unsaturated. How can you get rid of this smell?
Tasks for finding and explaining cause-and-effect relationships also play an important role in shaping students' ideas about the methods of scientific research, since causality is one of the forms of the general interconnection of the phenomena of the objective world. For many schoolchildren, completing assignments to determine the consequences of a theory is a rather difficult, but accessible type of work. No wonder scientists note that "the power of science is not only that it explains the observed phenomena, but also that it can predict the course of a process." Therefore, the essence of such tasks is questions like: “What is the reason for this?”, “How can this be explained?”, “Why did this happen?”, “What does it depend on?”, “What would change if ...?” Examples:
1. In a city where a plant for the production of phosphate fertilizers from fluorapatite concentrate began to operate, residents noticed that window panes were gradually fading. What are the possible reasons for this phenomenon?
2. What are the causes of acid rain? What impact do they have on: a) structures made of metal and concrete; b) technology; c) soil; d) works of art made of metal, marble, limestone?
One of the forms of implementation of the research method of teaching is the compilation of story-tasks, fairy tales, poetic works. This type of activity involves the writing by students of a small literary work that describes a phenomenon or substance veiled in the text. An archive of such student works has been collected in the office.
So do modern schoolchildren still need research skills? In my opinion, the words of the Nobel laureate, our countryman Zh. I. Alferov, whose opinion is certainly worthy of attention, can become an exhaustive answer to this question: “For every self-respecting country there are three privileged articles. In the first place, I put health care, because first of all a person must be physically healthy. On the second - education, because an uneducated person is not something that XXI century, but in the last century there was nothing to do. And I will put science in the third place, because it is science that determines the future of mankind ... ".
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In its pure form, oxygen was first obtained by Scheele in 1772, then in 1774 Priestley isolated it from mercury oxide.
The Latin name for oxygen "oxygenium" comes from the ancient Greek word "oxys", which means "sour", and "gennao" - "I give birth"; hence the Latin "oxygenium" means "acid-producing".
In the free state oxygen is found in air and water. In the air (atmosphere) it contains 20.9% by volume or 23.2% by weight; its content in water in the dissolved state is 7-10 mg/l.
In a bound form, oxygen is included in the composition of water (88.9%), various minerals (in the form of various oxygen compounds). Oxygen is part of the tissues of every plant. It is essential for the respiration of animals.
Oxygen occurs in nature in a free state in a mixture with other gases and in the form of compounds, and therefore both physical and chemical methods of obtaining it are used.
The general method for obtaining oxygen from compounds is based on the oxidation of a divalent negatively charged ion according to the scheme:
2O 2- - 4e - \u003d O 2.
Since the oxidation can be carried out in various ways, there are many different (laboratory and industrial) methods for obtaining oxygen.
1. DRY METHODS FOR OBTAINING OXYGEN BY THERMAL DISSOCIATION
Thermal dissociation of various substances can be carried out in test tubes, tubes, flasks and retorts made of refractory glass or in iron retorts.OBTAINING OXYGEN BY THERMAL DECOMPOSITION OF SOME METAL OXIDES (HgO, Ag 2 O, Au 2 O 3, IrO 2, etc.)
An experience. Thermal decomposition of red mercury oxide.2HgO \u003d 2Hg + O 2 - 2x25 kcal.
From 10 g of red mercury oxide, 500 ml of oxygen are obtained.
For the experiment, a test tube made of refractory glass 17 cm long and 1.5 cm in diameter with a 3-4 cm long lower end bent, as shown in , is used. 3-5 g of red mercury oxide is poured into the lower end. A rubber stopper with a drain tube is inserted into a test tube fixed in a support in an inclined position, through which the oxygen released during heating is removed into the crystallizer with water.
When red mercury oxide is heated to 500°, oxygen is released from the outlet tube and droplets of metallic mercury appear on the walls of the test tube.
Oxygen is poorly soluble in water, and therefore it is collected using the method of displacing water after complete removal of air from the device.
At the end of the experiment, the outlet tube is first removed from the crystallizer with water, then the burner is extinguished and, taking into account the toxicity of mercury vapor, the stopper is opened only after the tube has completely cooled.
Instead of a test tube, you can use a retort with a receiver for mercury.
An experience. Thermal decomposition of silver oxide. Reaction equation:
2Ag 2 O \u003d 4Ag + O 2 - 13 kcal.
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-oxide-sample.jpg) |
When black powder of silver oxide is heated in a test tube with an outlet tube, oxygen is released, which is collected over water, and a shiny layer of silver remains on the walls of the test tube in the form of a mirror.
PRODUCTION OF OXYGEN BY THERMAL DECOMPOSITION OF OXIDES WHICH, WHEN RECOVERED, TRANSITION TO OXIDES OF LOWER VALENCE, RELEASE PART OF OXYGEN
An experience. Thermal decomposition of lead oxides. As a result of intermolecular redox reactions, oxygen is released:A) 2PbO 2 \u003d 2PbO + O 2;
b) 2Pb 3 O 4 \u003d 6PbO + O 2;
PbO2 290-320°→ Pb 2 O 3 390-420°→ Pb 3 O 4 530-550°→PbO.
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 Red lead (Pb 3 O 4 or 2PbO PbO 2) |
 Red lead |
 Lead oxide (IV) PbO 2 |
 Lead oxide (IV) PbO 2 |
During thermal decomposition, about 460 ml of oxygen are obtained from 10 g of lead dioxide, and about 160 ml of oxygen from 10 g of Pb 3 O 4.
Obtaining oxygen from lead oxides requires more intense heating.
With strong heating of the dark brown powder PbO 2 or orange Pb 3 O 4, a yellow powder of lead oxide PbO is formed in the test tube; with the help of a smoldering splinter, you can make sure that oxygen is being released.
The test tube after this experiment is not suitable for further use, since. when heated strongly, lead oxide combines with glass.
An experience. Thermal decomposition of manganese dioxide.
3MnO 2 \u003d Mn 3 O 4 + O 2 - 48 kcal.
About 420 ml of oxygen are obtained from 10 g of manganese dioxide (pyrolusite). In this case, the test tube is heated to a light red heat.
To obtain a large amount of oxygen, the process of decomposition of pyrolusite is carried out in an iron tube 20 cm long closed at one end. The second end is closed with a stopper with a tube through which oxygen is removed.
The iron tube is heated using an incineration furnace or a Teklu gas burner with a dovetail nozzle.
An experience. Thermal decomposition of chromic anhydride. Oxygen is formed as a result of an intramolecular redox reaction:
4CrO 3 \u003d 2Cr 2 O 3 + 3O 2 - 12.2 kcal.
 Chromium (VI) oxide CrO 3 [chromic anhydride] |
 Chromium oxide (III) Cr 2 O 3 |
 Chromium oxide (III) Cr 2 O 3 |
Thermal decomposition of chromic anhydride (a hygroscopic, dark red solid) releases oxygen and forms a green chromium oxide powder, Cr 2 O 3 .
OXYGEN PRODUCTION BY THERMAL DECOMPOSITION OF PEROXIDES
An experience. Thermal decomposition of barium peroxide BaO 2 . The reversible reaction proceeds as follows:2ВаO 2 + 38 kcal ← 500° 700°→ 2ВаО + O 2 .
With strong heating of barium peroxide BaO 2, the peroxide bond is broken with the formation of barium oxide and the release of oxygen.
About 660 ml of oxygen is obtained from 10 g of barium peroxide.
Instead of barium peroxide, sodium peroxide can also be used. Then the expansion goes according to the equation
2Na 2 O 2 \u003d 2Na 2 O + O 2.
The experiment is carried out in a test tube with an outlet tube.
An experience. Thermal decomposition of potassium chlorate. Potassium chlorate decomposes differently depending on the temperature. When it is heated to 356°, it melts, and at 400° it decomposes according to the equation
2KSlO 3 \u003d KClO 4 + KCl + O 2.
In this case, only one third of the oxygen contained in the compound is released and the melt solidifies. This phenomenon is explained by the fact that the resulting compound KClO 4 is more stable and refractory.
When potassium chlorate is heated to 500°, the formation of potassium perchlorate is an intermediate reaction. The expansion in this case proceeds according to the equations:
A) 4KSlO 3 = 3KSlO 4 + KCl + 71 kcal;
b) 3KSlO 4 = 3KSl + 6O 2 - 24 kcal;
4KSlO 3 \u003d 4KSl + 6O 2 + 52 kcal.
Thermal decomposition of potassium chlorate is carried out in a small retort, which is connected by means of a drain tube with a safety tube to a crystallizer filled with water (or a pneumatic bath). The device is assembled in accordance with. To avoid an explosion, pure KClO 3 is poured into the retort, without admixture of organic substances.
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To avoid violent decomposition, which may cause the retort to burst, heating is carried out carefully.
The released oxygen is collected in various vessels above the water. When they want to get a slow flow of oxygen, potassium chlorate is diluted by mixing it with dry table salt.
An experience. Thermal decomposition of potassium chlorate in the presence of a catalyst. In the presence of catalysts (MnO 2 , Fe 2 O 3 , Cr 2 O 3 and CuO), potassium chlorate easily and completely decomposes at a lower temperature (without the formation of an intermediate compound, potassium perchlorate) according to the equation:
2KSlO 3 \u003d 2KSl + 3O 2 + 19.6 kcal.
When manganese dioxide is added, KClO 3 decomposes already at 150-200 °; The process has the following intermediate steps:
2KSlO 3 + 6MnO 2 → 2KSl + 6MnO 3 → 2KSl + 6MnO 2 + 3O 2 + 19.6 kcal.
The proportion of added manganese dioxide (pyrolusite) is from 5 to 100% by weight of potassium chlorate.
The test tube with potassium chlorate is closed with a stopper, through which two glass tubes are passed. One tube serves to remove oxygen into the crystallizer with water, the second, a very short tube, bent at a right angle with a closed outer end, contains a fine powder of black manganese dioxide MnO 2 .
The device is assembled in accordance with. When the test tube is heated to approximately 200°, no oxygen bubbles are released in the crystallizer with water. But as soon as you turn up the short tube with manganese dioxide and lightly knock on it, a small amount of manganese dioxide will enter the test tube and a rapid evolution of oxygen will immediately begin.
After the end of the experiment and the cooling of the apparatus, the mixture of manganese dioxide and potassium chloride is poured into water. After the dissolution of potassium chloride, the hardly soluble manganese dioxide is filtered off, washed thoroughly on the filter, dried in an oven and stored for further use as a catalyst. If it is necessary to obtain a large amount of oxygen, the decomposition process is carried out in refractory glass retorts or in cast iron retorts.
Thermal decomposition of potassium chlorate in the presence of manganese dioxide is the most convenient of the dry methods for obtaining oxygen.
This experiment is done with other catalysts - Fe 2 O 3 , Cr 2 O 3 and CuO.
An experience. Obtaining oxygen by heating potassium chlorate, a mixture of potassium chlorate with manganese dioxide and manganese dioxide. The following instruments are required for the experiment: three test tubes made of refractory glass with outlet tubes, three cylinders with a capacity of 100 ml each, three gas burners, three crystallizers and three racks with clamps.
The installation is assembled in accordance with. The crystallizers and cylinders are filled with water, slightly tinted with potassium permanganate or fuchsin S.
Pour 1 g of pure KClO 3 into the first tube, 0.5 g of KClO 3 and 0.5 g of MnO 2 into the second, and 1 g of MnO 2 into the third. Particular attention is paid to ensure that the test tubes are clean and that no cork grains get into them.
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Carefully adjusted gas burners, burning with the same, not very strong, non-luminous flame and emitting the same amount of heat, are placed under the test tubes so that they heat the substance in the test tube with the top of the flame.
Soon, oxygen begins to be released from the tube containing the mixture of potassium chlorate and manganese dioxide, and the reaction ends before it begins to be released in other tubes.
Increasing the heating of the remaining two test tubes. As soon as the potassium chlorate melts and oxygen begins to be released, reduce the flame so that there is no violent outgassing. In a test tube with manganese dioxide, oxygen begins to be released only after the contents of the test tube are heated to red heat. Oxygen liberated from each tube is collected in crystallizers by displacing colored water from the cylinders.
At the end of the experiment, the burners are extinguished, the outlet tubes are removed, then manganese dioxide is isolated from the middle test tube in the manner described above.
The conducted experiment clearly shows the features of these three different methods of obtaining oxygen.
OXYGEN PRODUCTION BY THERMAL DECOMPOSITION OF BROMATES AND IODATES
The behavior of these salts during heating was considered when studying the properties of bromates and iodates. Their decomposition is carried out in test tubes with outlet tubes; the oxygen released is collected over water.OBTAINING OXYGEN BY THERMAL DECOMPOSITION OF NITRATE
According to how nitrates decompose when heated, they can be divided into three groups:1. Nitrates decomposing as a result of intramolecular redox reactions to nitrites and oxygen. This group includes alkali metal nitrates. The reactions proceed according to the equations:
2NaNO 3 \u003d 2NaNO 2 + O 2,
2KNO 3 \u003d 2KNO 2 + O 2.
2. Nitrates decomposing as a result of intramolecular redox reactions into metal oxide, nitrogen dioxide and oxygen. This group includes nitrates of all metals, with the exception of alkali and noble metals. For example:
2Pb (NO 3) 2 \u003d 2PbO + 4NO 2 + O 2,
2Cu(NO 3) 2 \u003d 2CuO + 4NO 2 + O 2,
2Hg (NO 3) 2 \u003d 2HgO + 4NO 2 + O 2.
3. Nitrates decomposing as a result of intramolecular redox reactions into metal, nitrogen dioxide and oxygen. This group includes noble metal nitrates:
2AgNO 3 \u003d 2Ag + 2NO 2 + O 2.
The unequal decomposition of nitrates during heating is explained by the different stability of the corresponding nitrites and oxides.
Alkali metal nitrites are stable, lead (or copper) nitrites are unstable, but their oxides are stable, and as for silver, both nitrites and oxides are unstable; therefore, when nitrates of this group are heated, free metals are released.
An experience. Thermal decomposition of sodium or potassium nitrate. Sodium or potassium nitrate is heated in a test tube or retort with an outlet tube. At 314° sodium nitrate melts, and at 339° potassium nitrate; only after the contents in the test tube or retort become red-hot, does the decomposition of nitrate begin according to the equations given above.
Decomposition proceeds much more easily if the melting of nitrates is prevented by mixing them with manganese dioxide or soda lime, which is a mixture of NaOH and CaO.
The thermal decomposition of lead and silver nitrates is considered in experiments on the production of nitrogen dioxide.
OXYGEN PRODUCTION BY THERMAL DECOMPOSITION OF PERMANGANATES
An experience. Thermal decomposition of potassium permanganate. Reaction equation:2KMnO 4 \u003d K 2 MnO 4 + MnO 2 + O 2.
This intramolecular redox reaction occurs at approximately 240°. Thermal decomposition is carried out in a dry test tube (or retort) with a gas outlet tube. If you want to get pure oxygen without traces of dust, which is formed during thermal decomposition, a glass wool swab is inserted into the neck of the test tube (or retort).
This is a convenient way to get oxygen, but it is expensive.
After the end of the experiment and cooling of the test tube (or retort), several milliliters of water are poured into it, the contents are thoroughly shaken and the color of the formed substances is observed (K 2 MnO 4 is green and MnO 2 is dark brown).
Due to the property of potassium permanganate to release oxygen when heated, it is used along with sulfur, coal and phosphorus in various explosive mixtures.
 Obtaining oxygen by thermal decomposition of potassium permanganate |
 Na 2 MnO 4 |
 Manganese dioxide МnO 2 |
 Manganese dioxide МnO 2 |
 Manganese dioxide МnO 2 |
OXYGEN PRODUCTION BY THERMAL DECOMPOSITION OF PERSULPHATES
An experience. For the experiment, freshly prepared ammonium persulfate is used, since it changes its composition during storage. Ammonium persulfate (solid) decomposes on heating according to the following equation:(NH 4) 2 S 2 O 8 \u003d (NH 4) 2 SO 4 + SO 2 + O 2.
To free oxygen from sulfur dioxide impurities, the gas mixture is passed through a NaOH solution, which binds sulfur dioxide in the form of sodium sulfite. Thermal decomposition is carried out in a test tube with an outlet tube.
OXYGEN PRODUCTION BY THERMAL DECOMPOSITION OF PERCHLORATES
This method is considered when describing the experience of obtaining oxygen by thermal decomposition of potassium chlorate without a catalyst; in this case perchlorate is the intermediate.OXYGEN PRODUCTION BY THERMAL DECOMPOSITION OF PERCARBONATES
An experience. Sodium percarbonate, when heated, decomposes according to the equation:2K 2 C 2 O 6 \u003d 2K 2 CO 3 + 2CO 2 + O 2.
To free oxygen from carbon dioxide impurities, the gas mixture is passed through a solution of calcium or barium hydroxide.
Oxygen can also be obtained by burning oxygenite. Oxygenite is called a thin mixture of 100 wt. including KClO 3 , 15 wt. including MnO 2 and a small amount of coal dust.
The oxygen obtained by this method is contaminated with an admixture of carbon dioxide.
Along with substances that decompose with the release of oxygen when heated, there are many substances that do not release oxygen when heated. To verify this, they do experiments with heating CuO, CaO, Na 2 SO 4, etc.
II. WET METHODS OF OBTAINING OXYGEN
OXYGEN PRODUCTION BY DECOMPOSITION OF ALKALI METAL PEROXIDES WITH WATER
The reaction proceeds according to the equation:2Na 2 O 2 + 4H 2 O \u003d 4NaOH + 2H 2 O + O 2.
This is a highly exothermic reaction that proceeds in the cold and is accelerated by catalysts - salts of copper, nickel, cobalt (for example, CuSO 4 .5H 2 O, NiSO 4 .7H 2 O and CoSO 4 .7H 2 O).
Convenient for obtaining oxygen is oxylite - a mixture of sodium peroxide Na 2 O 2, potassium K 2 O 2 and anhydrous copper sulfate. This mixture is stored in tightly closed iron boxes, protected from atmospheric moisture (which decomposes it, see the equation of the previous reaction) and carbon dioxide, with which it reacts according to the equation:
Na 2 O 2 + 2СO 2 = 2Na 2 СO 3 + O 2 + 113 kcal.
An experience. A pinch of sodium peroxide (or oxylite) is poured into a test tube (glass or flask) with a small amount of cold water; in this case, a rapid release of oxygen is observed and the vessel is heated.
If the experiment is carried out in a vessel with an outlet tube, then the released oxygen can be collected.
PRODUCTION OF OXYGEN BY DECOMPOSITION OF PEROXIDES WITH ACIDS IN THE PRESENCE OF CATALYSTS, FOR EXAMPLE MnO 2 OR PbO 2
An experience. Add dilute HCl to a test tube with barium peroxide and manganese dioxide; in this case, oxygen is released as a result of the reaction:2ВаO 2 + 4НCl = 2ВаСl 2 + 2Н 2 O + O 2.
When PbO 2 is used as a catalyst, dilute HNO 3 is added to the mixture.
OXYGEN PRODUCTION BY CATALYTIC DECOMPOSITION OF HYDROGEN PEROXIDE
Reaction equation:2H 2 O 2 \u003d 2H 2 O + O 2.
When studying the properties of hydrogen peroxide, factors favoring its decomposition are noted, and experiments are carried out on its decomposition under the influence of manganese dioxide and colloidal silver solution.
An experience. In a glass cylinder with 50 ml of water and 10-15 ml of perhydrol(30% solution of H 2 O 2) add some finely divided powder of manganese dioxide; there is a rapid release of oxygen with the formation of foam (this phenomenon is very similar to boiling).
The experiment can also be done in a test tube, and instead of perhydrol, a 3% hydrogen peroxide solution can be used.
Instead of MnO 2, you can use a colloidal solution of silver.
PRODUCTION OF OXYGEN BY THE ACTION OF POTASSIUM PERMANGANATE ON HYDROGEN PEROXIDE (IN ACID, NEUTRAL AND ALKALINE ENVIRONMENTS)
The reaction proceeds according to the equations below; hydrogen peroxide is the reducing agent:2KMnO 4 + 3H 2 SO 4 + 5H 2 O 2 \u003d 2MnSO 4 + K 2 SO 4 + 8H 2 O + 5O 2,
2KMnO 4 + 2H 2 O + 3H 2 O 2 \u003d 2MnO 2 + 2KOH + 4H 2 O + 3O 2,
2KMnO 4 + 2KOH + H 2 O 2 \u003d 2K 2 MnO 4 + 2H 2 O + O 2.
An experience. Obtaining an easily regulated direct current of oxygen by oxidizing hydrogen peroxide in the cold potassium permanganate in an alkaline medium. A 3-5% solution of hydrogen peroxide acidified with a 15% solution of H 2 SO 4 is poured into a Bunsen flask, and a 10% solution of potassium permanganate is poured into a dropping funnel fixed in the neck of the flask.
With the help of a dropping funnel tap, both the flow of the permanganate solution into the flask and the flow of oxygen can be controlled. During the experiment, a solution of KMnO 4 is introduced into the flask drop by drop.
The Bunsen flask in the experiment can be replaced by a Wurtz flask or a two-necked flask.
An experience. Production of oxygen by oxidation of hydrogen peroxide with manganese dioxide in an acidic medium. Reaction equation:
MnO 2 + H 2 SO 4 + H 2 O 2 \u003d MnSO 4 + 2H 2 O + O 2.
The reaction proceeds in the cold; therefore, for the experiment, you can use any device that allows the interaction in the cold between a solid and a liquid substance to obtain a constant flow of gas (Kipp apparatus or a Wurtz flask, a Bunsen flask or a two-necked flask with a dropping funnel).
During the experiment, manganese dioxide in pieces, 15% H 2 SO 4 and 3-5% hydrogen peroxide solution are used.
An experience. Obtaining oxygen by oxidation of hydrogen peroxide with potassium ferricyanide in an alkaline medium. Reaction equation:
2K 3 + H 2 O 2 + 2KOH \u003d 2K 4 + 2H 2 O + O 2.
The reaction proceeds in the cold; to obtain a direct current of oxygen, the devices indicated in the previous experiment are used, solid potassium ferricyanide, a 6-10% solution of potassium oxide hydrate and a 3-5% solution of hydrogen peroxide.
An experience. Obtaining oxygen by heating chromate (bichromate or chromic anhydride) with concentrated sulfuric acid. Due to the reversible reaction proceeding according to the equation:
2CrO 4 2- + 2H + ↔ Cr 2 O 7 2- + H 2 O,
in an acidic environment, dichromate is always present, not chromate.
The following reactions take place between concentrated sulfuric acid and bichromate:
K 2 Cr 2 O 7 + H 2 SO 4 \u003d 2CrO 3 + K 2 SO 4 + H 2 O,
(reaction of double exchange and dehydration)
4CrO 3 + 6H 2 SO 4 \u003d 2Cr 2 (SO 4) 3 + 6H 2 O + 3O 2.
(redox reaction)
When conducting an experiment in a test tube, oxygen is released and the orange color (characteristic of dichromate) changes to green (characteristic of trivalent chromium salts).
III. OBTAINING OXYGEN FROM LIQUID AIR
To liquefy air, the principle is used, according to which, when a gas expands without external work, a significant decrease in temperature occurs (the Joule-Thomson effect).Most gases heat up when compressed and cool when expanded. A schematic diagram of the operation of a Linde machine used to liquefy air is shown.
Compressor B with the help of a piston compresses up to 200 atm the air supplied through valve A, purified from carbon dioxide, moisture and traces of dust. The heat generated during compression is absorbed in cooler D cooled by running water. After that, valve C is opened and air enters vessel E, where it expands to a pressure of 20 atm. Due to this expansion, the air is cooled to approximately -30°. From vessel E, the air returns to compressor B; passing through the outer tube of the coil G, it cools along the way a new portion of compressed air, which goes towards it along the inner tube of the coil. The second portion of air is thus cooled to approximately -60°. This process is repeated until the air is cooled to -180°; this temperature is sufficient to liquefy it at 20 atm in vessel E. The liquid air accumulating in vessel E is drained into a cylinder through valve 1. The described installation operates continuously. The details of this machine are not shown in the diagram. This machine was improved by J. Claude, after which it became more productive.
In its composition, liquid air differs from ordinary atmospheric air; it contains 54% by weight liquid oxygen, 44% nitrogen, and 2% argon.
An experience. To show how the properties of organic substances change under the influence of changing conditions (temperature and oxygen concentration), plants with leaves and flowers or a thin rubber tube are immersed in a thermos with liquid air using metal tongs.
Oxygen is obtained from liquid air in the following ways:
- a) fractional distillation (the most common method);
- b) dissolving air in liquids (for example, 33% oxygen and 67% nitrogen dissolve in water) and extracting it under vacuum;
- c) selective absorption (charcoal absorbs 92.5% by volume of oxygen and 7.5% by volume of nitrogen);
- d) based on the difference in the diffusion rates of oxygen and nitrogen through the rubber membrane.
To purify oxygen, it is passed through a wash bottle with alkali, which retains all acidic volatile compounds accompanying it, through a solution of KI (to free it from ozone) and through concentrated H 2 SO 4, which retains water vapor.
PROPERTIES OF OXYGEN
PHYSICAL PROPERTIES
Oxygen is a colorless, odorless and tasteless gas.Its density relative to air is 1.10563; therefore, it can be collected in vessels using the air displacement method.
Under normal conditions, one liter of oxygen weighs 1.43 g, and one liter of air weighs 1.29 g. The boiling point is -183°, the melting point is -218.88°.
Liquid oxygen in a thin layer is colorless, thick layers are blue; the specific gravity of liquid oxygen is 1.134.
Solid oxygen is blue and looks like snow; its specific gravity is 1.426.
The critical temperature of oxygen is -118°; critical pressure 49.7 atm. (Oxygen is stored in steel cylinders with a capacity of 50 liters, at a pressure of 150 atm. Methods for storing various gases in steel cylinders are described in the first chapter.)
In water, oxygen dissolves in a very small amount: in one liter of water at 20 ° and a pressure of 760 mm Hg. Art. 31.1 ml of oxygen dissolves. Therefore, it can be collected in test tubes, cylinders or gasometers using the water displacement method. Oxygen dissolves better in alcohol than in water.
To use a gasometer (), you must be able to fill it with water and gas under atmospheric pressure, as well as above and below atmospheric pressure; be able to release gas from the gasometer.
First, gasometer A is filled with water through funnel B, with taps C and D open and hole E closed. Water, entering the gasometer from funnel B through tap C, displaces air from it through tap D.
To fill the gasometer with gas at a certain pressure, close valve C and B and open hole E: if both upper valves are tightly fitted, water does not flow out of the gasometer. The end of the tube is inserted through hole E, through which gas flows under pressure exceeding atmospheric pressure. Gas accumulates in the upper part of the gasometer, displacing water from it, which pours out through hole E. After the gas is almost completely filled with gasometer, hole E is closed. When filling the gasometer with gas under atmospheric or reduced pressure, the tube through which the gas enters is connected to an open valve B, then hole E is opened and valve C is left closed. Water flowing out of hole E sucks the gas into the gasometer. After the gasometer is almost completely filled with gas, close hole E and valve B.
To release gas, fill funnel B with water, open tap C; Water, entering the gasometer, displaces gas from it, which exits through the open tap E).
In the molten state, some metals, such as platinum, gold, mercury, iridium and silver, dissolve about 22 volumes of oxygen, which is released when they solidify with a specific sound, especially characteristic of silver.
The oxygen molecule is very stable, it consists of two atoms; at 3000° only 0.85% of the oxygen molecules dissociate into atoms.
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 Gasometers are not only laboratory. The photo shows the Vienna Gasometers - these are 4 large structures located in Vienna (Austria) and built in 1896-1899. They are located in Simmering, the eleventh district of the city. In 1969-1978, the city abandoned the use of coke oven gas in favor of natural gas, and the gas meters were closed. In 1999-2001 they were rebuilt and became multifunctional complexes (Wikipedia).
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CHEMICAL PROPERTIES
According to its chemical activity, oxygen is second only to fluorine.It combines with other elements directly or forms compounds indirectly. The direct connection of oxygen can proceed vigorously and slowly. The combination of oxygen with elements or complex substances is called oxidation or combustion. It always proceeds with the release of heat, and sometimes light. The temperature at which oxidation occurs can vary. Some elements combine with oxygen in the cold, others only when heated.
In the event that during a chemical reaction the amount of heat released exceeds its losses as a result of radiation, thermal conductivity, etc., vigorous oxidation occurs (for example, combustion of metals and non-metals in oxygen), otherwise slow oxidation occurs (for example, phosphorus, coal, iron, animal tissue, pyrite, etc.).
If the slow oxidation proceeds without heat loss, there is a rise in temperature, which leads to an acceleration of the reaction, and the slow reaction can become vigorous as a result of self-acceleration.
An experience. An example of self-acceleration of a slow reaction. Take two small pieces of white phosphorus. One of them is wrapped with filter paper. After a while, a piece of phosphorus wrapped in paper ignites, while an unwrapped one continues to slowly oxidize.
There is no clear line between vigorous and slow oxidation. Vigorous oxidation is accompanied by the release of a large amount of heat and light; slow oxidation is sometimes accompanied by cold luminescence.
Combustion also proceeds differently. Substances that, during combustion, turn into a vapor state (sodium, phosphorus, sulfur, etc.), burn with the formation of a flame; substances that do not form gases and vapors during combustion burn without a flame; the combustion of some metals (calcium, magnesium, thorium, etc.) is accompanied by the release of a large amount of heat, and the hot oxides formed in this case have the ability to emit a lot of light in the visible region of the spectrum.
Substances that release a large amount of heat during oxidation (calcium, magnesium, aluminum) are able to displace other metals from their oxides (aluminothermy is based on this property).
Combustion in pure oxygen is much more vigorous than in air, in which it slows down due to the fact that it contains about 80% nitrogen, which does not support combustion.
COMBUSTION OF VARIOUS SUBSTANCES IN OXYGEN
Experiments illustrating combustion in oxygen are carried out in thick-walled and wide-mouthed flasks with a capacity of 2.5-3 l (), on the bottom of which a thin layer of sand should be poured (if this is not done, then when a drop of molten metal hits the bottom of the vessel, the vessel may burst ).For combustion in oxygen, the substance is placed in a special spoon made of a thick iron (or copper) wire flattened at the end, or a sample to be burned is attached to the end of the wire.
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An experience. Ignition and combustion in oxygen of a smoldering splinter (or candle). When a smoldering splinter (or candle) is introduced into a vessel with oxygen, the splinter ignites and burns with a bright flame. Sometimes a splinter ignites with a small explosion. The experience described is always used to discover free oxygen ( * Nitrous oxide gives a similar reaction).
An experience. Combustion of coal in oxygen. Reaction equation:
C + O 2 \u003d CO 2 + 94.3 kcal.
If a piece of smoldering coal, fixed at the end of an iron wire, is introduced into a vessel with oxygen, the coal burns out with the release of a large amount of heat and light. The carbon dioxide formed during combustion is discovered using blue litmus paper moistened with water or by passing the gaseous products of combustion through a solution of calcium hydroxide.
The experience of burning coal in oxygen released during the thermal decomposition of KClO 3 has already been carried out in the study of the properties of potassium chlorate.
An experience. Burning sulfur in oxygen. Reaction equation:
S + O 2 \u003d SO 2 + 71 kcal.
When an ignited sulfur color is introduced into a vessel with oxygen, a more intense burning of sulfur in oxygen is observed and a sharp smell of sulfur dioxide is felt. To prevent this poisonous gas from spreading throughout the laboratory, the vessel is tightly closed at the end of the experiment.
The combustion of sulfur in oxygen released during the thermal decomposition of potassium chlorate was described in the study of the properties of KClO 3 .
An experience. Combustion of white and red phosphorus in oxygen. The reaction proceeds according to the equation:
4P + 5O 2 \u003d 2P 2 O 5 + 2x358.4 kcal.
The short and wide neck of a flask (or jar) with a capacity of 0.5-2 liters, placed on a tray of sand, is closed with a cork with a metal spoon and a glass tube passed through it, the axis of which should pass through the middle of the spoon ().
Simultaneously with filling the flask with oxygen (by displacing air), a pea-sized piece of white phosphorus is cut off in a mortar under water, lightly compressed with filter paper to remove traces of water, and placed in a metal spoon with metal tongs. The spoon is lowered into the flask, it is closed, and the phosphorus is touched with a glass rod (or wire) heated to 60-80°C, which is inserted through a glass tube.
Phosphorus ignites and burns with a bright flame to form phosphorus pentoxide as white smoke (causing cough).
Sometimes white phosphorus ignites in oxygen without being touched by a heated glass rod or wire. Therefore, it is recommended to use phosphorus stored in very cold water; it should be squeezed out with filter paper without any friction, and in general, all preparations for introducing it into a vessel with oxygen should be carried out as quickly as possible. If phosphorus After the combustion of phosphorus, take out the cork with a spoon, pour a small amount of water into the flask and test it with blue litmus paper.
If part of the phosphorus remains unoxidized, the spoon is lowered into the crystallizer with water. If all the phosphorus has burned out, then the spoon is calcined under a draft, washed with water and dried over a burner flame.
In carrying out this experiment, molten white phosphorus is never introduced into the vessel with oxygen. This cannot be done, firstly, because phosphorus can be easily spilled, and, secondly, because in this case phosphorus burns in oxygen too violently, scattering splashes in all directions that can fall on the experimenter; splashes of phosphorus burst a vessel, fragments of which can injure others.
Therefore, there should be a crystallizer with water on the table, into which phosphorus can be thrown in case it catches fire when it is squeezed with filter paper; it is also necessary to have a concentrated solution of KMnO 4 or AgNO 3 (1: 10) for first aid in case of phosphorus burns.
Dry red phosphorus can be used instead of white phosphorus. To do this, red phosphorus is pre-purified, thoroughly washed with water and dried.
Red phosphorus ignites at a higher temperature, so it is set on fire with a very hot wire.
After burning, and in this case, pour a little water into the flask, test the resulting solution with litmus and ignite the spoon under draft.
In both experiments, goggles made of dark glass should be used.
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An experience. Combustion in oxygen of metallic sodium. The reaction proceeds according to the equation:
2Na + O 2 = Na 2 O 2 + 119.8 kcal.
Sodium is burned in a small crucible made of pure calcium oxide, chalk or asbestos cardboard, but not in a metal spoon, which, from the heat released when sodium is burned in oxygen, can itself melt and burn.
Sodium is set on fire and brought into a vessel with oxygen, in which it burns with a very bright flame; its combustion should be observed through protective dark glasses.
A crucible prepared from chalk (or CaO) is attached with two or three thin wires to a thick iron (or copper) wire () and a pea-sized piece of metallic sodium, cleaned of oxide, is placed in it.
Chalk, asbestos, calcium oxide are poor conductors of heat, and therefore ignite sodium by directing a burner flame at it from above with a blowpipe. To protect yourself from splashes of burning sodium, a rubber tube is put on the blowpipe.
Heating, melting and ignition of sodium in air is carried out over a vessel with oxygen.
If the sodium does not ignite, then a crust formed on the metal surface is removed with a blowpipe, but this should be done with extreme caution because of the possible splashing of molten sodium.
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An experience. Combustion in oxygen of metallic calcium. Reaction equation:
2Ca + O 2 \u003d 2CaO + 2x152.1 kcal.
A match is placed in a small crucible made of asbestos cardboard, and calcium chips are placed on top of it.
Light a match and bring the crucible with calcium chips into a vessel with oxygen. Through protective goggles, the ignition and combustion of calcium metal with a bright flame is observed.
Ignited calcium can also be added to a vessel with oxygen (as was done in the previous experiment with sodium).
An experience. Combustion of magnesium in oxygen. The reaction proceeds according to the equation:
2Mg + O 2 \u003d 2MgO + 2x143.84 kcal.
A piece of tinder is attached to one end of a magnesium tape 20-25 cm long, twisted in the form of a spiral, and an iron wire is attached to the other. The wire is taken in hand and, holding the magnesium ribbon in a vertical position, tinder is set on fire and the magnesium ribbon is introduced into a vessel with oxygen. Magnesium ignites and burns through goggles to form magnesium oxide.
At the end of the experiment, a little water is poured into the vessel and, with the help of an indicator, they are convinced of the alkaline nature of the solution of the formed magnesium hydroxide.
The experiment can be done with magnesium powder. To do this, take a spoonful of magnesium powder and insert half a match with a head into it. Light a match and put a spoon into a vessel with oxygen.
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However, magnesium burns with a dazzling flame in air, although here the oxidative reactions of oxygen are significantly weakened due to the fact that the air contains a large percentage of nitrogen.
A vessel in which magnesium is burned can burst if burning magnesium is not introduced into it quickly enough or if burning magnesium touches its walls.
The bright light of burning magnesium has found application for illuminating photographed objects, and also as an initiator of some reactions occurring under the influence of short light waves, for example, the synthesis of HCl from elements.
When considering the properties of potassium chlorate, the experience of burning its mixture with magnesium was described.
An experience. Combustion in oxygen of large zinc filings. Reaction equation:
2Zn + O 2 \u003d 2ZnO + 2x83.17 kcal.
Large zinc filings are poured into a refractory glass tube 15 cm long and an inner diameter of 0.8-1 cm (in the absence of them, powder can also be used, but in such a way that oxygen can pass through it) and strengthen it at one end in a horizontal position in the tripod clamp.
The end of the tube fixed in a tripod is connected to a source of oxygen, and the opposite end is heated with a gas burner.
When oxygen is passed through a tube, zinc ignites and burns with a bright flame to form zinc oxide (a white solid). The experiment is carried out under pressure.
An experience. Determination of the amount of oxygen consumed during the combustion of copper.
2Cu + O 2 \u003d 2CuO + 2x37.1 kcal.
The device for the experiment is shown in. A porcelain boat with 1 g of fine powder of metallic copper is inserted into a refractory tube 20 cm long and 1.5 cm in inner diameter. The wash bottle with water is connected to an oxygen source (gasometer or cylinder).
The gasometer with a bell, located on the right, is filled with water, tinted with a solution of indigo or magenta. The gasometer valve is opened so that the oxygen passing through the device can flow under the bell.
Open the clamp between the wash bottle and the refractory tube and let in about 250 ml of oxygen under the bell. Close the clamp and note the exact volume of oxygen.
With the help of a Teklu dovetail burner, the part of the tube in which the porcelain boat is located is heated. After a few minutes, the copper lights up and the water level in the bell immediately rises.
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Heating is continued for 35-40 minutes until the volume of gas in the gasometer stops changing.
Allow the device to cool down. this sets a constant volume of gas. Then the water is brought to the same level and the volume of unreacted oxygen is determined from the divisions of the gasometer.
The experiment makes it possible to accurately determine the amount of oxygen consumed for the oxidation of copper weighed before the start of the experiment.
Do not use this appliance to burn zinc, magnesium or calcium powder.
An experience. Confirmation of the law of constancy of composition. Accurately, to hundredths of a gram, an empty porcelain crucible with a lid is weighed, which had previously been thoroughly cleaned, calcined and cooled in a desiccator. Then approximately 3-4 g of fine copper powder is poured into the crucible and the crucible with copper is accurately weighed.
Place the crucible in an inclined position on a porcelain triangle and heat it over low heat for 15-20 minutes. The lid is then removed and heated strongly with an oxidizing burner flame. After 20-25 minutes, cover the crucible with a lid and continue heating. After stopping the heating, the crucible is cooled in a desiccator and accurately weighed.
- g 1 = weight of an empty crucible with a lid;
- g 2 = weight of the empty crucible with lid and copper;
- g 3 = weight of the empty crucible with lid and copper oxide.
By repeating the experiment with metallic copper and other metals, they find that in all cases oxygen combines with various elements in a constant quantitative ratio, and in practice they are convinced that the ratio between the weight of the substances entering into a chemical compound is always constant.
An experience. Combustion of iron in oxygen. Reaction equation:
4Fe + 3O 2 \u003d 2Fe 2 O 3 + 2x196.5 kcal.
For the experiment, a thin wire made of tempered steel with a diameter of 7-8 mm is used, one end of which is stuck into a cork stopper, and a piece of tinder is attached to its other end or wrapped with thread and immersed in molten sulfur (sulfur wick). When a steel spiral with a lit tinder (or sulfur wick) is introduced into a vessel with oxygen (at the bottom of which there should be a layer of sand), the spiral burns out, scattering sparks.
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An experience. Combustion of metal powders in air. Above the flame of a gas burner installed under the draft, a pinch of powder of copper, zinc, iron, magnesium, aluminum, antimony is poured.
An experience. Oxidation of metals in a closed vessel. Experience allows us to prove that during the transformation of metals into oxides, part of the air is consumed and that the increase in the weight of metals during their oxidation is equal to the weight loss of air.
A test tube with fine iron powder is tightly closed with a rubber stopper, through which a glass tube must be passed with a rubber tube put on it with a screw clamp (). The stopper and clamp should seal the tube hermetically.
After weighing the assembled apparatus, the test tube is heated with a flame of a gas burner with continuous shaking until sparks form in the powder. After the tube has cooled, it is checked by weighing on a balance whether the weight of the tube has changed. Then a glass tube is inserted into the rubber tube, the end of which is lowered into a glass of water.
When opening the clamp, watch how the water rises through the tube. This is due to the fact that the oxygen in the air was used up for the oxidation of iron and therefore the pressure in the device decreased.
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To detect a small difference between the weight of iron and the weight of iron oxide is possible only with the help of sufficiently sensitive balances.
Instead of a test tube, you can use a retort or a round-bottomed flask, and instead of a rubber stopper, you can use a waxed cork stopper.
Similar experiments were carried out by Lomonosov and Lavoisier to prove the law of conservation of matter.
An experience. Slow oxidation of wet iron. Experience allows us to establish that heat is released during the oxidation of wet iron powder.
The device consists of a thermoscope connected to a pressure gauge (). Two tubes are introduced into the reaction space of the thermoscope through a tightly fitted rubber stopper. The first tube is connected to a gas cylinder and serves to supply oxygen. The second tube serves to remove the gas; it is connected to a Müncke washing bottle, into which water is poured, tinted with indigo or magenta.
Such an amount of water is poured into the washing flask that, when sucked into the inner tube and filled, there is still water left in the flask, which would close the outlet of the tube.
For the manufacture of a thermoscope, you can use the outer part of the Drexel 300 ml wash bottle with a side tube. A test tube 23 cm long and 2.5 cm in diameter with a slightly narrowed neck is inserted into the vessel. The upper outer part of the tube should be ground to the neck of the vessel. In the absence of the above parts, the thermoscope can be made from a Bunsen flask, into the neck of which a large test tube is inserted using a rubber ring. The thermoscope is connected to a U-shaped pressure gauge, into which water tinted with magenta is poured.
The pressure gauge has a T-tap with a stopcock, which facilitates its adjustment.
In a conical flask, 100 g of iron powder are mixed with benzene, filtered through a folded filter, washed with ether and quickly (oxidized iron powder is not suitable for experiment) dried on a tile of porous ceramic material.
Iron powder, thoroughly moistened with 18 ml of distilled water, is scattered over glass wool and filled with it throughout the reaction space of the thermoscope.
To remove air from the device, a strong jet of oxygen is blown through it. The presence of pure oxygen in the apparatus is determined by bringing a smoldering splinter to the outlet of the wash bottle. Then stop the flow of oxygen and equalize the liquid in both tubes of the manometer (behind the manometer is strengthened graph paper).
In the reaction vessel, oxygen is partially combined with iron, and after a few minutes, absorption of liquid into the inner tube of the wash bottle is observed. In this case, some more oxygen is passed into the thermoscope to equalize the liquid levels in the inner and outer tubes of the wash bottle. This operation is repeated two or three times. The change in pressure indicated by the manometer indicates the release of heat during oxidation.
The section on phosphorus describes experiments showing the slow oxidation of white phosphorus.
An experience. Catalytic oxidation of methyl alcohol to formaldehyde. The reaction proceeds according to the equation:
H 3 C-OH + 0.5O 2 → H 2 C \u003d O + H 2 O + 36 kcal.
The device is assembled in accordance with. 50 ml of pure methyl alcohol is poured into a 150 ml Wurtz flask with the end of the side tube extended to a diameter of 1 mm. In a refractory tube 25-30 cm long and 1 cm in diameter, a roll of copper mesh 10 cm long wound on a thick copper wire is inserted. Water is poured into the washing flask on the left, and a colorless solution of sulfurous acid H 2 SO 3 with fuchsin is poured into the flask on the right just before the start of the experiment. The glass into which the Wurtz flask is lowered must contain water heated to 30-40 °.
To conduct the experiment, water is heated in a glass to 45-48 °, a strong current of air is sucked through the device using a water jet pump, and a copper grid roller is heated with a Teklu burner, first with a weak flame, then brought to red heat.
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The air current is regulated in such a way that after the burner is removed, the copper grid roller remains red-hot without heating from the outside.
After some time, the mixture of sulfurous acid with fuchsin in the right wash bottle turns into an intense red-violet color.
In parallel, it is shown that the reaction of a formaldehyde solution with a colorless solution of sulfurous acid and fuchsin is characteristic of aldehyde.
To obtain a colorless solution of sulfurous acid with fuchsin, 0.1 g of fuchsin is dissolved in 300 ml of distilled water and sulfur dioxide is passed through the resulting solution until the color of fuchsin disappears. The resulting reagent is stored in a vial with a ground stopper. The entire experience lasts about five minutes. At the end of the experiment, allow the apparatus to cool in a weak current of air.
When using ethyl alcohol, acetaldehyde is formed according to the equation:
CH 3 CH 2 -OH + 0.5O 2 → CH 3 CH \u003d O + H 2 O.
The reduction of an oxidized roll from a copper grid with methyl alcohol is described in the section on nitrogen (a method for obtaining nitrogen by binding atmospheric oxygen with hot copper).
An experience. Anodic oxidation, bleaching effect of oxygen at the moment of its release. A glass with a solution of sodium sulfate is covered with a cork circle, through which two carbon electrodes with a diameter of 5-6 mm are passed.
The anode is wrapped several times with a blue-dyed cotton cloth and the electrodes are connected to three batteries connected in series.
After 2-3 minutes of current passing, the first two layers of tissue, directly adjacent to the anode, become discolored by atomic oxygen released during electrolysis. The second and subsequent layers of tissue, through which already stable diatomic oxygen molecules pass, remain colored.
An experience. anodic oxidation. A 25% solution of H 2 SO 4 is poured into a glass and two lead electrodes in the form of plates are lowered into it. The electrodes are connected to a source of direct electric current with a voltage of 10 V. When the circuit is closed, a brown color appears at the anode.
The electrolysis is continued until the brown lead dioxide PbO 2 formed on the anode becomes visible.
If you use a silver anode, then black oxide of silver Ag 2 O is released on the anode.
Putting out the fire. Knowing what combustion is, it is easy to understand what fire extinguishing is based on.
Fire can be extinguished with solids, gases and vapours, liquids and foams. To extinguish the fire, it is necessary to isolate it from air (oxygen), for which it is thrown with sand, salt, earth or covered with a thick blanket.
Fire extinguishers are often used to extinguish fires, which are described in the section on carbon dioxide.
When extinguishing burning wood warehouses, straw, textiles, paper, so-called dry fire extinguishers are used, which emit solid carbon dioxide having a temperature of -80 ° C. In this case, the flame goes out due to a strong decrease in temperature and dilution of the oxygen in the air with carbon dioxide, which does not support combustion. These fire extinguishers are handy for fires in power plants, telephone exchanges, oil and varnish factories, distilleries, etc.
An example of the use of gases for extinguishing fires is the use of sulfur dioxide, which is formed during the combustion of sulfur thrown into a furnace or chimney, to extinguish soot that has ignited in a furnace chimney.
The most common and cheapest fire extinguishing fluid is water. It lowers the temperature of the flame, and its vapors prevent air from reaching burning objects. However, water is not used to extinguish burning oil, gasoline, benzene, oil and other flammable liquids lighter than water, as they float to the surface of the water and continue to burn; the use of water in this case would only contribute to the spread of fire.
Foaming fire extinguishers are used to extinguish gasoline and oils; the foam they throw out remains on the surface of the liquid and isolates it from the oxygen in the air.
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OXYGEN APPLICATIONS
Oxygen is used as an oxidizing agent in the production of nitric, sulfuric and acetic acids, in the blast furnace process, for underground coal gasification, for gas welding and cutting of metals (hydrogen or acetylene-oxygen flame), for melting metals, quartz, to obtain high temperatures in laboratories, for breathing with the use of various devices used by pilots, divers and firefighters.Without oxygen, no animal can exist.
Coal, oil, paraffin, naphthalene and a number of other substances impregnated with liquid oxygen are used to prepare some explosives.
Mixtures of liquid oxygen with coal powder, wood flour, oil and other combustible substances are called oxyliquites. They have very strong explosive properties and are used in demolition work.
OZONE O 3
Ozone is an allotropic form of oxygen. The name comes from the Greek word "osein", which means "smelly". Ozone was discovered in 1840 by Shenbein.Ozone is found in very small quantities in the atmosphere: at the earth's surface, its concentration is 10 -7%, and at a height of 22 km from the earth's surface - 10 -6%. On the surface of the earth, ozone is found mainly near waterfalls, on the seashore (where it, like atomic oxygen, is formed under the influence of ultraviolet rays), in coniferous forests (here it is formed as a result of the oxidation of terpenes and other organic substances); ozone is formed during lightning discharges. At an altitude of about 22 km from the earth's surface, it is formed from oxygen under the influence of the ultraviolet rays of the sun.
Ozone is produced from oxygen; in this case, it is necessary to expend external energy (thermal, electrical, radiation). The reaction proceeds according to the equation:
3O 2 + 69 kcal ↔ 2O 3.
Thus, the conversion of oxygen to ozone is an endothermic reaction in which the volume of gases decreases.
Oxygen molecules under the influence of thermal, light or electrical energy break down into atoms. Being more reactive than molecules, atoms enter into combination with undissociated oxygen molecules and form ozone.
The amount of ozone formed is the greater, the lower the temperature, and almost does not depend on the pressure at which the reaction proceeds. It is limited by the decay rates of the resulting ozone molecules and their formation as a result of photochemical action (during electrical discharges, under the influence of radiation from quartz lamps).
With all methods of obtaining ozone under conditions close to ordinary temperature, its low yield (about 15%) is characteristic, due to the instability of this compound.
The decomposition of ozone can be partial (when it proceeds spontaneously at ordinary temperature; in this case it is proportional to the concentration) and complete (in the presence of catalysts).
 The stratosphere at an altitude of 15-35 km contains the ozone layer, which protects the Earth from ultraviolet radiation. Many have heard of the so-called "ozone hole". In reality, this is only a partial decrease in the ozone content, which is significant only over the south pole of the planet. But even here, the destruction of the ozone layer is only partial. It is possible that the "ozone hole" was formed long before the emergence of mankind. Significant amounts of ozone are also formed near the surface of the planet. One of the main sources is anthropogenic pollution (especially in big cities). This ozone is far from harmless - it poses a significant risk to human health and the environment. - Ed. |
 Distribution of ozone over the southern hemisphere September 21-30, 2006 Blue, purple and red indicate areas with low ozone content, green and yellow areas with higher ozone content. NASA data. (ed. note) |
CHEMICAL METHODS FOR PRODUCING OZONE
All oxygen production reactions result in the formation of small amounts of ozone.An experience. Production of ozone by the action of concentrated sulfuric acid on potassium permanganate. Reaction equations:
- 2KMnO 4 + H 2 SO 4 \u003d 2HMnO 4 + K 2 SO 4 (exchange reaction),
- 2HMnO 4 + H 2 SO 4 \u003d Mn 2 O 7 + H 2 O + H 2 SO 4 (dehydration reaction),
- Mn 2 O 7 → 2MnO 2 + 3O,
- Mn 2 O 7 → 2MnO + 5O (both redox decomposition reactions can occur simultaneously; more vigorous decomposition leads to the formation of MnO),
- 3O + 3O 2 = 3O 3 (ozone formation reaction).
The manganese anhydride Mn 2 O 7 formed according to the above equations is a heavy greenish-brown oily liquid that decomposes at 40-50 ° into MnO 2, MnO and atomic oxygen, which, when combined with molecular oxygen in the air, forms ozone.
Instead of a mortar, you can use a porcelain cup, watch glass or asbestos tiles.
Introduced into the atmosphere of ozone at the tip of the wire, a lump of cotton wool soaked in ether immediately ignites. Instead of ether, cotton wool can be moistened with alcohol, gasoline or turpentine.
Water-moistened starch-iodide indicator paper turns blue with ozone. This phenomenon is explained by the reaction:
2KI + O 3 + H 2 O \u003d I 2 + 2KOH + O 2.
Starch iodine paper is obtained by wetting strips of filter paper in a mixture of a colorless concentrated solution of potassium iodide and a starch solution.
The blue color of the starch iodine paper gradually disappears as the reaction proceeds between iodine and potassium hydroxide:
3I 2 + 6KOH = KIO 3 + 5KI + 3H 2 O.
In the presence of excess ozone, free iodine is oxidized; the following reactions take place:
I 2 + 5O 3 + H 2 O \u003d 2HIO 3 + 5O 2,
I 2 + 9O 3 \u003d I (IO 3) 3 + 9O 2.
 Interaction of Mn 2 O 7 with wool |
An experience. Production of ozone by the action of concentrated nitric acid on ammonium persulfate. The source of atomic oxygen in this experiment is persulfuric acid, which is formed as a result of the exchange reaction between ammonium persulfate and nitric acid, and the source of molecular oxygen is nitric acid decomposing when heated.
This method of producing ozone is based on the following reactions:
(NH 4) 2 S 2 O 8 + 2HNO 3 \u003d H 2 S 2 O 8 + 2NH 4 NO 3, 
2HNO 3 → 2NO 2 + 0.5O 2 + H 2 O,
O + O 2 \u003d O 3.
The device necessary for the experiment is shown in. A small flask containing 2 g of ammonium persulfate and 10 ml of concentrated nitric acid is connected by means of a thin section to a glass tube, the end of which is lowered into a test tube with a solution of potassium iodide and a small amount of starch.
Some time after the start of heating the flask on low heat, the solution in the test tube turns blue. However, as a result of the interaction of iodine with potassium hydroxide, the blue color soon disappears.
A 0.5% solution of indigo carmine or a 1% solution of indigo in concentrated H 2 SO 4 changes color from blue to pale yellow due to ozone oxidation of indigo to isatin according to the equation:
C 16 H 10 O 2 N 2 + 2O 3 ← 2C 8 H 5 O 2 N + 2O 2 + 63.2 kcal.
Instead of a cone in this experiment, you can use a test tube with a gas outlet tube.
White phosphorus, previously cleaned from the surface film under water, is placed using metal tongs in a glass cylinder with a capacity of 1.5-2 liters.
Distilled water is poured into the cylinder so that it covers 2/3 of the phosphorus sticks, and it is placed in a crystallizer with water heated to 25 °.
The cylinder can be replaced by a 500 ml flask, in which the phosphorus can be heated until it melts (approximately 44 °) with continuous agitation.
The presence of ozone is detected approximately two hours after the start of the experiment by a characteristic smell reminiscent of garlic and indicator starch iodine paper; ozone can be detected by pouring a few drops of titanyl sulfate into a test tube with a solution taken from the cylinder.
Titanyl sulfate is obtained by heating under draft in a porcelain cup 1 g of titanium dioxide with a double volume of concentrated sulfuric acid until white fumes are released. After cooling, the contents of the cup are gradually introduced into 250 ml of ice water. In water, titanium sulfate Ті (SO 4) 2 turns into titanyl sulfate.
In the presence of ozone, a colorless solution of titanyl sulfate turns into a yellow-orange solution of pertitanic acid, the reaction proceeds according to the equation:
TiOSO 4 + O 3 + 2H 2 O \u003d H 2 TiO 4 + O 2 + H 2 SO 4.
PRODUCTION OF OZONE BY ELECTROLYSIS OF ACIDS
An experience. Obtaining ozone by electrolysis of concentrated (approximately 50%) sulfuric acid. During the electrolysis of concentrated H 2 SO 4, redox processes on the electrodes proceed according to the following scheme:- H 2 SO 4 → HSO 4 - + H + (ions of concentrated sulfuric acid),
- H 2 O ↔ OH - + H + (water ions),
- At the cathode: 2H + 2e - → 2H → H 2 (hydrogen is released),
- At the anode: HSO 4 - - 2e - → H 2 S 2 O 8.
- Persulfuric acid decomposes in water according to the equation: H 2 S 2 O 8 + 2H 2 O \u003d 2H 2 SO 4 + H 2 O + O (oxygen is released at the anode).
O + O 2 \u003d O 3.
Depending on the conditions (current density and temperature), persulfuric acid, ozone, and molecular oxygen are formed at the anode.
During the electrolysis of acidified water, ozone is formed when the anode is made of a non-oxidizing metal, and the water does not contain substances capable of absorbing oxygen.
The device is assembled in accordance with. 100 ml of a 20-50% sulfuric acid solution is poured into a glass with a capacity of 150 ml, into which a cathode made of a lead plate (25 x 10 mm) and an anode, which is a platinum wire with a diameter of 0.5 mm, soldered into a glass plate, are immersed. tube 9 cm long and 5 mm in diameter. The wire is soldered in such a way that its free end protrudes 1 cm from the tube. The platinum wire is connected to the external wire with a few drops of mercury introduced into the tube. The anode is inserted through a waxed cork plug into an open tube 9 cm long and 1.5 cm in diameter, which has a side tube in the upper part.
After closing the electrical circuit, at a current strength of 1.5 A, ozone can be detected at the opening of the side tube by smell or using starch iodine paper.
If a platinum anode is used and the cell is cooled down to -14°, ozone can also be obtained in a small amount by electrolysis of dilute H 2 SO 4 .
Ozone is also obtained by electrolysis of chromic, acetic, phosphoric and hydrofluoric acids.
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PRODUCTION OF OZONE WITH ELECTRIC DISCHARGE IN OXYGEN
An experience. Obtaining ozone by passing electrical sparks through the oxygen contained in the eudiometer. In a Bunsen eudiometer (see the section on hydrogen) with platinum electrodes with a capacity of 50 ml, filled with a solution of potassium iodide containing starch, 5 ml of oxygen is introduced. The eudiometer is fixed with a tripod in the crystallizer with the same solution.When the wires of the eudiometer are connected to the secondary terminals of the induction coil, sparks jump between the platinum wires and the starched solution of potassium iodide begins to turn blue. Oxidation of the iodide solution by ozone is enhanced by shaking it.
Instead of the Bunsen eudiometer, you can use the device indicated on, made of thick glass. This device could ozonate all the introduced oxygen if there was no heating from spark discharges, which accelerates the reverse reaction of ozone decomposition.
A solution of potassium iodide with the addition of starch is prepared as follows: 0.5 g of starch is ground in a mortar in a small amount of water, the resulting dough is introduced with stirring into 100 ml of boiling water; after the starch solution has cooled, 0.5 g of KI, previously dissolved in a small amount of water, is added to it.
When a current of pure and dry oxygen (air) is passed through the ozonator under the action of a quiet electric discharge of electric discharges without sparks), some of the oxygen (maximum 12-15% by volume) is converted into ozone.
Humid and dusty air cannot be used for this purpose, since during electrical discharges in this case a thick fog is formed, which settles on the electrodes and glass walls of the ozonizer; as a result, instead of quiet discharges, sparks begin to jump in the ozonator, and nitric oxide is formed; nitric oxide in the presence of oxygen is oxidized to nitrogen dioxide, which destroys the electrodes.
The source of oxygen can be a gasometer or an oxygen cylinder; oxygen entering the ozonator is first passed through a wash bottle with concentrated H 2 SO 4 .
Under the action of such electric discharges in the space occupied by oxygen, ions and electrons are formed, which, when colliding with oxygen molecules, cause their decay.
The presence of ozone is detected by the methods described above, as well as by the methods indicated in the description of the properties of ozone.
Below are descriptions of some types of ozonizers.
By introducing alternately a layer of glass wool with manganese or lead dioxide powder (10 cm) or a layer of activated granular carbon into a wide tube, one is convinced that ozone decomposes when passing through them.
The decomposition of ozone is accompanied by the release of heat and an increase in the volume of the gas.
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OZONE APPLICATIONS
As a strong oxidizing agent, ozone kills microorganisms and is therefore used to disinfect water and air, to bleach straw, feathers, as an oxidizing agent in organic chemistry, in the production of ozonides, and also as a means of accelerating the aging of cognacs and wines.HYDROGEN PEROXIDE H 2 O 2
Hydrogen peroxide was first obtained in 1818 by Tenard by reacting barium peroxide with hydrochloric acid. |
SPREAD
In the free state, H 2 O 2 is found in the lower layers of the atmosphere, in precipitation (during lightning discharges, about 11 mg per 60 kg of water), as a product of the slow oxidation of organic and inorganic substances, as an intermediate product of assimilation and dissimilation, and in the juices of some plants.RECEIVING
An experience. Preparation of hydrogen peroxide by cathodic reduction of molecular oxygen with hydrogen. The reaction proceeds according to the equation:O 2 + 2H → H 2 O 2 + 138 kcal.
The device is assembled in accordance with. The electrolytic bath is a glass with a capacity of 250-300 ml, filled with sulfuric acid (sp. weight 1.2-1.25) and covered with an asbestos plate.
An anode and a glass cylinder 3 cm in diameter are passed through the plate, inside of which there is a cathode, as well as a glass tube through which pure oxygen is supplied from a gasometer or cylinder. An oxygen supply tube with a retracted tip passes from below under the cylinder and ends at the cathode itself.
Near the anode, another hole is made in the asbestos plate to remove oxygen released from the anode.
The anode is a platinum plate located at a higher level than the cathode. The cathode is made from a platinum or palladium plate.
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source electrical energy is a 10 V battery.
After assembling the device, 10 ml of electrolyte is taken from the anode space with a pipette, poured into a beaker, and a few drops of titanyl sulfate solution are added. No staining occurs in this case.
5-10 minutes after the start of electrolysis, with a current of 4-5 A and a strong jet of oxygen, the current is turned off and an electrolyte sample is taken. This time, when titanyl sulfate is added, the electrolyte turns yellow-orange; this is due to the formation of peroxodisulfatotitanic acid:
With a longer electrolysis, the samples for titanyl sulfate give a more intense color. In this case, the following reactions take place:
A) TiOSO 4 + H 2 O 2 + H 2 O \u003d H 2 TiO 4 + H 2 SO 4,
b) TiOSO 4 + H 2 O 2 + H 2 SO 4 \u003d H 2 [TiO 2 (SO 4) 2] + H 2 O.
An experience. Preparation of hydrogen peroxide by the action of dilute acids on alkaline peroxides (Na 2 O 2 or K 2 O 2). The reaction proceeds according to the equations:
Na 2 O 2 + H 2 SO 4 \u003d H 2 O 2 + Na 2 SO 4,
K 2 O 2 + H 2 SO 4 \u003d H 2 O 2 + K 2 SO 4.
The experiment is carried out in a test tube. Obtaining this method of hydrogen peroxide is not very convenient because of the difficulty of its separation from alkaline sulfates.
It is also impossible to recommend the production of hydrogen peroxide by the action of water on alkaline peroxides, since in these reactions hydrogen peroxide is only an intermediate compound, which decomposes into oxygen and water in the presence of alkalis; therefore, the reaction of interaction between alkaline peroxides and water underlies one of the wet methods for producing oxygen.
An experience. Obtaining hydrogen peroxide from barium peroxide by the action of sulfuric acid. Reaction equation:
BaO 2 + H 2 SO 4 \u003d H 2 O 2 + BaSO 4.
120 ml of water is poured into a glass, 5 ml of concentrated H 2 SO 4 (sp. weight 1.84) are added and it is immersed in a crystallizer with a mixture of ice and salt. Putting some ice in a glass at 0°C, gradually with continuous stirring add a suspension of barium peroxide, which is obtained by grinding in a mortar 15 g BaO 2 with 30 ml of ice water. The suspension is a hydrate of barium peroxide BaO 2 8H 2 O.
After filtering off the barium sulfate, a 3-5% hydrogen peroxide solution is obtained. A slight excess of acid does not interfere with the production of peroxide.
The presence of hydrogen peroxide is discovered as follows: pour 2 ml of the test solution and 2 ml of H 2 SO 4 into a test tube, add ether (layer 0.5 cm thick) and add a few drops of potassium chromate solution. In the presence of hydrogen peroxide in an acidic medium, chromates (as well as dichromates) form intensely colored perchromic acids, and the reaction proceeds:
H 2 Cr 2 O 7 + 4H 2 O 2 \u003d 2H 2 CrO 6 + 3H 2 O.
Perchromic acid H 2 CrO 6 with the structural formula
It is colored blue and decomposes already at room temperature; so the color of the solution quickly disappears. The ether extracts the acid from the solution when shaken and makes it more stable.
Peroxide compounds of chromium are reduced to compounds of trivalent chromium (green) with the release of oxygen.
An experience. Hydrogen peroxide can also be obtained from the hydrolysis of sodium perborate and barium percarbonate. In this case, the reaction proceeds according to the equations:
NaBO 3 + H 2 O \u003d NaBO 2 + H 2 O 2,
ВаС 2 O 6 + Н 2 O \u003d ВаСО 3 + CO 2 + Н 2 O 2.
PROPERTIES OF HYDROGEN PEROXIDE
Under normal conditions, hydrogen peroxide is a colorless, odorless liquid with an unpleasant metallic taste.At maximum concentration, it is a syrupy liquid with a specific gravity of 1.5. In a thick layer it has a blue color.
It dissolves in water, ethyl alcohol, ethyl ether in any ratio. On sale, hydrogen peroxide is usually found in the form of a 3% and 30% solution in distilled water. The latter is called "perhydrol". Under pressure 26 mm Hg. Art. boils at 69.7°. Hardens at -2°.
More stable are dilute solutions of hydrogen peroxide; as for concentrated solutions, they decompose with an explosion according to the equation:
2H 2 O 2 \u003d 2H 2 O + O 2 + 47 kcal.
The decomposition of hydrogen peroxide is favored by light, heat, certain inorganic and organic substances, glass roughness, and traces of dust.
From inorganic substances, hydrogen peroxide decomposes oxides (MnO 2, Fe 2 O 3, Cr 2 O 3), alkaline hydrates of NaOH, KOH, Ba (OH) 2 oxides in the presence of impurities, hydrated salts of Cu 2+, Co 3+, Pb ions 2+, Mn 2+, etc., ions of trivalent metals Fe 3+, Al 3+, metals in a highly crushed, especially in a colloidal state (Au, Ag, Pt), silicon compounds, including those that are part of the glass .
Organic substances that decompose hydrogen peroxide include blood, which activates decomposition due to the enzyme catalase contained in it, while its other enzyme, peroxidase, promotes the elimination of oxygen peroxide in the presence of oxidizing substances.
The catalytic decomposition of H 2 O 2 in the presence of alkalis, manganese dioxide and colloidal silver solution is described in the section “Production of oxygen by wet methods”.
An experience. Decomposition of hydrogen peroxide under the influence of heat. A flask with a capacity of 200-250 ml is filled almost completely with a solution of hydrogen peroxide; close with a stopper with a gas outlet tube, the tip of which is lowered into a crystallizer with water (). After removing air from the device, the flask is heated and the released oxygen is collected in a cylinder filled with water.
The flow of oxygen is regulated by increasing or decreasing the heating of the flask.
The presence of oxygen is discovered with a smoldering splinter.
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An experience. Catalytic decomposition of hydrogen peroxide. Approximately the same amount of perhydrol (30% hydrogen peroxide solution) is poured into three glasses. Manganese dioxide is added to the first glass, platinum black to the second, and a few drops of blood to the third.
Decomposition proceeds best in the third glass, where blood has been added. If sodium cyanide is added to the blood, and then perhydrol, oxygen is released weakly.
It has been experimentally established that colloidal platinum and catalase are poisoned by the same substances, for example, HCN, KCN, NaCN, CO, I 2 , H 2 S, CS 2, etc. The poisoning of catalysts is explained by the fact that their large surface adsorbs a significant amount of toxic substances . In this case, poisonous substances isolate the active surface of the catalyst from the reacting substance, and the catalyst loses its ability to accelerate the reaction.
An experience. Catalytic decomposition of hydrogen peroxide in an alkaline medium. To obtain glowing in the dark water, four solutions are prepared:
- 1) dissolve 1 g of pyrogallol C 6 H 3 (OH) 3 powder in 10 ml of distilled water;
- 2) dissolve 5 g of K 2 CO 3 in the same amount of distilled water;
- 3) take 10 ml of 35-40% solution of formaldehyde CH 2 O;
- 4) take 15 ml of a 30% solution of hydrogen peroxide (perhydrol).
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When the eyes adjust to the darkness, pour the perhydrol into the glass with continuous stirring. The liquid begins to boil, as it were, foams and glows with a yellow-orange light, shimmering with brilliant foam.
The release of light during chemical reactions that occur without appreciable release of heat is called chemiluminescence. The light emitted by chemiluminescence is most often red or yellow. In the present experiment, chemiluminescence is explained by the oxidation of pyrogallol with hydrogen peroxide in an alkaline medium. The energy released during oxidation is almost entirely converted into light, although a small amount of it is also released in the form of thermal energy, which heats the contents of the glass and causes partial evaporation of formaldehyde (a pungent odor spreads).
Instead of pyrogallol, hydroquinone, resorcinol, or photographic developers can be used.
Hydrogen peroxide can be made more stable by adding to it a small amount of one of the following substances (stabilizers): barbituric acid, uric acid, phosphoric acid, sulfuric acid, sodium phosphate, urea, phenacetin, etc.
Hydrogen peroxide is a very weak acid (weaker than carbonic acid). Its acidic properties can be determined using a neutral solution of litmus.
Two types of salts correspond to hydrogen peroxide: hydroperoxides (NaHO 2, KHO 2) and peroxides (Na 2 O 2, K 2 O 2, BaO 2).
In chemical reactions, hydrogen peroxide can act both as an oxidizing agent and as a reducing agent.
Sometimes a very slight change in pH leads to a radical change in the redox function of hydrogen peroxide. The following reactions are examples:
I 2 + 5H 2 O 2 → 2HIO 3 + 4H 2 O; at pH1 H 2 O 2 oxidizing agent,
2НIO 3 + 5Н 2 O 2 → I 2 + 6Н 2 O + 5O 2; at pH2 H 2 O 2 reducing agent.
As an oxidizing agent, hydrogen peroxide breaks down as follows:
H-O-O-H → H-O-H + O.
(the liberated oxygen atoms react with the reducing agent, turning into negatively charged divalent oxygen).
OXIDATION WITH HYDROGEN PEROXIDE IN ACID MEDIUM
Oxidation of a negatively charged iodine ion with hydrogen peroxide is described in the section on obtaining free iodine. (This reaction is used to determine traces of hydrogen peroxide.)An experience. Oxidation of ferrous ion with hydrogen peroxide to ferric ion. Reaction equation:
2FeSO 4 + H 2 SO 4 + H 2 O 2 = Fe 2 (SO 4) 3 + 2H 2 O.
 FeSO4 |
 Fe 2 (SO 4) 3 |
In a test tube with a freshly prepared green solution of FeSO 4, dilute sulfuric acid and a 3% hydrogen peroxide solution are poured. Due to the oxidation of the divalent iron ion to the trivalent color of the solution changes and becomes yellow. The presence of the ferric ion can be determined using the thiocyanate ion, since ferric thiocyanate is intensely colored blood red (the reaction is very sensitive).
An experience. Hydrogen peroxide oxidation of sulfurous acid (sulfites) in sulfuric acid(sulfates). The reaction proceeds according to the equation:
H 2 SO 3 + H 2 O 2 \u003d H 2 SO 4 + H 2 O.
If hydrogen peroxide is added to an aqueous solution of sulfur dioxide (sulphurous acid), then sulfurous acid is oxidized to sulfuric acid.
In order to verify the formation of sulfuric acid, you can use the fact that BaSO 3 is soluble in mineral acids, while BaSO 4 is slightly soluble in them.
An experience. Oxidation of potassium ferricyanide with hydrogen peroxide. Reaction equation:
2K 4 + H 2 O 2 + H 2 SO 4 \u003d 2K 3 + 2H 2 O + K 2 SO 4.
If a slightly diluted H 2 SO 4 and a 3% solution of H 2 O 2 are added to a test tube with a yellow solution of potassium ferrocyanide, then the solution in the test tube turns brown-red, characteristic of potassium iron cyanide.
An experience. Oxidation of lead sulfide with hydrogen peroxide. The reaction proceeds according to the equation:
PbS + 4H 2 O 2 \u003d PbSO 4 + 4H 2 O.
To a solution of Rb(NO 3) 2 [or Rb(CH 3 COO) 2 ] add an aqueous solution of hydrogen sulfide; a black precipitate of lead sulfide precipitates. The reaction goes according to the equation:
Pb (NO 3) 2 + H 2 S \u003d PbS + 2HNO 3.
A 3% hydrogen peroxide solution is added to the precipitate of lead sulfide, thoroughly washed by decantation; oxidized to lead sulfate, the precipitate becomes white.
This reaction is based on the renewal of paintings blackened by time (due to the formation of lead sulfide on them).
An experience. Oxidation of indigo with hydrogen peroxide. When boiling in a test tube 5-6 ml of a dilute solution of indigo and 10-12 ml of a 3% or stronger solution of hydrogen peroxide, a discoloration of the indigo solution is observed.
OXIDATION WITH HYDROGEN PEROXIDE IN ALKALINE MEDIUM
An experience. Oxidation of chromites to chromates with hydrogen peroxide. The reaction proceeds according to the equation:2KCrO 2 + 2KOH + 3H 2 O 2 \u003d 2K 2 CrO 4 + 4H 2 O.
Hydrogen peroxide is added to a green solution of alkali metal chromite; chromite is oxidized to chromate and the solution turns yellow.
Alkali metal chromite is obtained by the action of alkali (in excess) on a solution of a trivalent chromium compound (see oxidation with bromine water in an alkaline medium).
An experience. Oxidation of bivalent manganese salts with hydrogen peroxide. Reaction equation:
MnSO 4 + 2NaOH + H 2 O 2 \u003d H 2 MnO 3 + Na 2 SO 4 + H 2 O.
Alkali is added to a colorless (or slightly pink) solution of a divalent manganese compound. A white precipitate of manganese hydroxide precipitates, which, even in the presence of traces of oxygen, oxidizes to manganese dioxide hydrate, and the precipitate turns brown.
Nitrous oxide in the presence of manganese dioxide hydrate forms manganese oxide.
The reactions described above proceed as follows:
MnSO 4 + 2NaOH \u003d Mn (OH) 2 + Na 2 SO 4,
Mn (OH) 2 + 1 / 2O 2 \u003d H 2 MnO 3 or MnO (OH) 2, 
In the presence of hydrogen peroxide, the oxidation of nitrous oxide to manganese dioxide hydrate proceeds very rapidly.
When heated, the oxidation of divalent manganese salts with hydrogen peroxide proceeds to the formation of manganese dioxide according to the equation:
MnSO 4 + H 2 O 2 + 2KOH = MnO 2 + K 2 SO 4 + 2H 2 O.
In a number of reactions, hydrogen peroxide serves as a reducing agent in both alkaline and acidic environments.
As a reducing agent, hydrogen peroxide decomposes as follows:
H-O-O-H → 2H + O=O.
Since peroxides can be both oxidizing and reducing agents, peroxide electrons can move from one molecule to another:
H 2 O 2 + H 2 O 2 \u003d O 2 + 2H 2 O.
Hydrogen peroxide reduction of KMnO 4 and MnO 2 in an acidic medium and K 3 in an alkaline medium is described in the wet oxygen production section.
An experience. Reduction of dark brown silver oxide to metallic silver with hydrogen peroxide. The reaction proceeds according to the equation:
Ag 2 O + H 2 O 2 \u003d 2Ag + H 2 O + O 2.
Pour into a test tube 2 ml of a dilute AgNO 3 solution, 4-6 ml of a 3% H 2 O 2 solution and 2-3 ml of a dilute NaOH solution. A black precipitate of metallic silver is formed according to the overall reaction equation:
2AgNO 3 + 2NaOH + H 2 O 2 \u003d 2Ag + 2NaNO 3 + 2H 2 O + O 2.
Under the action of alkalis on solutions of silver salts, instead of an unstable silver oxide hydrate, a dark brown precipitate of silver oxide precipitates (this property is also characteristic of hydrates of oxides of other noble metals).
In excess of alkalis, silver oxide is insoluble.
An experience. Recovery of gold compounds with hydrogen peroxide. Recovery can proceed both in acidic and alkaline environments.
In a test tube with a small amount of gold chloride solution, add a little alkali solution and a 3% hydrogen peroxide solution. There is an instant reduction of the trivalent gold ion to free gold:
2AuCl 3 + 3H 2 O 2 + 6KOH = 2Au + 6H 2 O + 3O 2 + 6KCl.
An experience. Hydrogen peroxide reduction of hypochlorites and hypobromites. Reaction equations:
KClO + H 2 O 2 \u003d KCl + H 2 O + O 2,
NaClO + H 2 O 2 \u003d NaCl + H 2 O + O 2,
NaBrO + H 2 O 2 \u003d NaBr + H 2 O + O 2,
CaOCl 2 + H 2 O 2 \u003d CaCl 2 + H 2 O + O 2.
These reactions form the basis of test-tube experiments for the production of oxygen.
Addition products of hydrogen peroxide. Such a substance is perhydrol - the product of the addition of hydrogen peroxide to urea:
This compound in the crystalline state is stabilized by traces of citric acid. When simply dissolved in water, hydrogen peroxide is formed.
Storage of hydrogen peroxide. Hydrogen peroxide is stored in a dark and cold place in paraffin (or glass waxed inside) vessels sealed with a paraffin stopper.
USE OF HYDROGEN PEROXIDE
A 3% solution of hydrogen peroxide is used in medicine as a disinfectant, for gargling and washing wounds; in industry they are used for bleaching straw, feathers, glue, ivory, furs, leather, textile fibers, wool, cotton, natural and rayon. A 60% solution is used to bleach fats and oils.Compared to chlorine, hydrogen peroxide has great advantages as a bleaching agent. It is used for the production of perborates (for example, sodium perborate, which is the active ingredient in bleach preparations).
Highly concentrated solutions of hydrogen peroxide (85-90%) mixed with some combustible substances are used to produce explosive mixtures.
WATER H 2 O
Cavendish was the first to synthesize water by burning hydrogen in 1781; its weight composition was accurately established by Lavoisier in 1783, and its volumetric composition - in 1805 by Gay-Lussac.SPREAD
Water is the most common hydrogen compound; it covers two-thirds of the earth's surface, filling the oceans, seas, lakes, rivers. Much water is in earth's crust, and in the form of vapors in the atmosphere.the purest natural water is the water of atmospheric precipitation, the most polluted with impurities is the water of the seas and oceans. By their nature, impurities can be inorganic and organic. In water, they can be in a dissolved and suspended state.
Water impurities are: free carbon dioxide, nitrogen, oxygen, CaCO 3 , Ca (HCO 3) 2 , MgCO 3 , CaSO 4 , MgSO 4 , alkali metal chlorides, silicic acid and its salts of alkali and alkaline earth metals, oxides of iron, aluminum , manganese, alkali and alkaline earth metal salts of nitric, nitrogenous and phosphoric acid, microorganisms and various organic substances in a colloidal state.
Mineral waters, in addition to these impurities, contain hydrogen sulfide, sulfates, salts of boric, arsenic, hydrofluoric, hydrobromic, hydroiodic and other acids.
An experience. Using the Ba 2+ ion, the presence of SO 4 2- ions is established in any natural water, using the Ag + ion, the presence of the Cl - ion, and by evaporating 500 ml of water in a cup, the presence of a dry residue.
RECEIVING
Getting water is described in the section on chemical properties hydrogen (hydrogen combustion). Water is formed when hydrogen combines with oxygen under the action of an electric discharge; the production of water is also described in the sections on the construction of eudiometers and the reduction of oxides with hydrogen.Water can be obtained by heating substances containing crystallization water, for example: CuSO 4 5H 2 O, Na 2 CO 3 10H 2 O, Na 2 B 4 O 7 10H 2 O, Na 2 SO 4 10H 2 O, FeSO 4 7H2O; as a by-product, it is formed during neutralization reactions, redox and other reactions.
To obtain large quantities of chemically pure water, none of the methods described above is used to obtain it, but resort to purification of very common natural water in various ways.
NATURAL WATER PURIFICATION
Physical impurities are separated by filtration through a regular or folded filter, a porous ceramic or glass plate, or through glass wool.To retain impurities that impart hardness to water, water is passed through permutite filters, and to get rid of coloring substances, through activated carbon.
The removal of impurities dissolved in water is achieved in the distillation process. The simplest distillation apparatus is shown, consisting of a Wurtz flask, a refrigerator and a receiver.
In order not to disassemble the device each time and to avoid connections with plugs, it is recommended to use a device made of Jena glass ().
Uniform boiling during distillation is achieved due to the fact that a little porous porcelain is first placed in the flask.
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The water obtained in this way contains gases in the dissolved state, for example CO 2, and a very small amount of silicates (formed as a result of the dissolution of the glass of the refrigerator by water condensate).
To remove gases (for example, CO 2), pour 750 ml of distilled water into a 1000 ml flask, throw a few pieces of capillary tubes into it and boil for 30-40 minutes. At the end of boiling, close the flask with a stopper, into which a tube with soda lime (a mixture of CaO and NaOH) is inserted. Soda lime absorbs carbon dioxide from the air, which can get into distilled water after it has cooled.
Since a large amount of distilled water is consumed in the chemical laboratory for the preparation of solutions and washing of precipitates, several continuous distillation apparatuses are described below.
Distillation apparatus Kaleshchinsky() consists of a retort with a side tube and a curved neck connected to a spiral cooler.
A constant water level in the retort and cooler is maintained by a siphon.
Before the start of the experiment, water is sucked into the siphon through the side tube, on which the rubber tube is to be put on, and the rubber tube is closed with a clamp or a glass rod is tightly inserted into it.
To ensure uniform boiling, several pieces of porous porcelain are placed in the retort before distillation begins, and a flask is attached to the end of the side tube of the siphon, into which air bubbles that enter the siphon when the water is heated will collect (air bubbles in the siphon can disrupt the normal supply of water to the retort) .
This small apparatus can operate continuously for quite a long time without requiring special care.
Verkhovsky distillation apparatus(). Description of the device: wide tube BUT serves to collect air bubbles released from the water when it is heated. She when filling the siphon B, C, D almost completely filled with water. Bottle F with a cut off bottom is closed with a cork with a tube passed through it E(to remove excess water from the bottle). All parts of the apparatus are interconnected by means of rubber plugs and tubes. Water from the faucet goes into the refrigerator, from there - into the bottle F, then - into the siphon B, C, D to the distillation flask. The same level of water in the flask and flask is maintained by means of a siphon B, C, D. The normal functioning of this, like the previous one, is ensured by the continuous flow of water from the tap.
In addition to those described, there are a number of other, more complex devices. Preference is given to devices made of Jena glass, in which the individual parts are connected not by plugs, but by sections. You can also use metal apparatus heated by electricity or gas.
Distilled water can be single, double and multiple distillation.
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PROPERTIES OF WATER
Water can be in solid, liquid and gaseous state. The transition from one state to another is determined by temperature and pressure.An experience. Difference between steam and fog. A small amount of water is poured into a 100 ml flask; a glass tube 5 cm long and 6 mm in diameter with a slightly drawn outer end is inserted into the neck of the flask. Having placed the flask on a tripod covered with an asbestos mesh, it is heated to an intense boil of water. The resulting water vapor is invisible both in the flask and at the tube opening, but clouds of fog (droplets of condensed vapor) form above the flask. For uniform boiling of water, several pieces of porous porcelain or glass beads are placed inside the flask.
It is not necessary to pull the end of the tube strongly, as this can create high pressure and then the flask will burst.
Clean water in all states of aggregation colorless. Water vapor is invisible.
An experience. Pairs, visible and invisible. Four large bottles are placed on the table. A little water is poured into the first, bromine into the second, alcohol into the third, and gasoline into the fourth.
After some time, the air in each flask is saturated with the vapors of the corresponding liquid. In a flask with bromine, vapors are visible; in flasks with water, alcohol and gasoline they are invisible; in bottles with alcohol and gasoline, they can be detected by smell.
The density of pure water at +4°C and a pressure of 760 mm Hg. Art. taken as a unit.
An experience. Confirmation that the density of warm water is less than that of water at +4°C. For the experiment, they use a glass tube bent in the form of a square, with each side about 25 cm long (). Both ends of the tube are connected with two pieces of rubber tube to a glass T-tube. The whole device is filled cold water, from which air must be removed by boiling, and fixed in a tripod in the position indicated in the figure. A few drops of ink, KMnO 4 solution, methylene blue or fluorescein are added to the T-tube and the dye is observed to diffuse in both directions. Then they heat the device at one of the corners and notice how the heated water, becoming lighter, begins to rise up and all the liquid in the tube begins to move in the direction indicated by the arrows in the figure. The dye from the T-shaped tube begins to move in the direction opposite to heating. If we now move the gas burner to the left corner, the colored water begins to move from left to right. This appliance serves as a central heating model.
Ice is less dense than water at +4°C, so it floats on liquid water.
An experience. Checking the weak thermal conductivity of water. Taking a test tube by the lower end, heat water in it. The water at the opening of the test tube begins to boil, remaining cold at its lower end, for which the test tube is held by hand.
The electrical conductivity of pure water is very low; pure water is a poor conductor of electricity.
An experience. To study the electrical conductivity of pure water and solutions of various electrolytes and non-electrolytes, a special device is used.
The main parts of the device for determining the electrical conductivity of liquids are: two electrodes, a lamp base with an electric lamp, a socket, a plug, a breaker, an electric current source and an electric wire.
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The electrodes can be platinum, carbon or copper; lamps can be of different power, but they prefer to use lamps used for flashlights; the current source can be 1-2 batteries or rectifiers, as well as transformers connected to the electrical network and giving a voltage of 3-4 V.
The electrodes are switched on with a plug. Instead of a base with an electric lamp, you can use an electric bell. Usually, the device (base with electric lamp, socket and breaker) is mounted on the same board according to the diagram shown in.
At the lower end of the electrodes, a mark is made to which it is necessary to pour liquid into the vessel when the electrodes are immersed in it.
copper electrodes. Two copper wires 10-12 cm long and 0.5-0.8 cm in diameter.
Both electrodes, like the previous ones, are fixed in a cork circle, into which a dropping funnel is also inserted.
To determine the electrical conductivity, liquid can be poured into a test tube, glass, cylinder, flask or jar, depending on the size of the electrodes used.
To conduct the experiment, the electrodes are immersed in a liquid and connected to an electrical circuit connected in series with an electric lamp (bell) and through a switch with a source of electrical energy.
If the light turns on (or the bell rings) when the current is turned on, then the liquid is a good conductor of electricity.
Every time before testing the electrical conductivity of a new liquid, the electrodes, the vessel into which the test liquid is poured, and the funnel are thoroughly washed with distilled water, alcohol, ether, chloroform, toluene or other solvent and wiped with filter paper.
Usually, the electrical conductivity of the following liquids is checked in the laboratory: distilled water, dilute solutions of HCl, H 2 SO 4, NaOH, Ba (OH) 2, NaCl and sugar.
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To show that electrical conductivity is due to the presence of ions, it is sufficient to demonstrate the following:
- a solution of Ba (OH) 2 + phenolphthalein conducts an electric current;
- H 2 SO 4 solution conducts electricity.
At atmospheric pressure (760 mm Hg), water boils at 100°. If the pressure changes, the boiling point of water also changes.
An experience. Boiling water at reduced pressure. The device is assembled in accordance with. It consists of a Liebig refrigerator with an inner tube made of thick and durable glass, ending at the bottom with a small cone. At the opposite end of the tube from the cone, there should be a hook for hanging the thermometer.
A little water is poured into the flask of the refrigerator, the thermometer is hung so that its ball with mercury is in the water of the flask, and the refrigerator is fixed in a vertical position on a tripod.
The inner tube of the refrigerator is connected through a safety vessel and a pressure gauge to a water jet pump.
At the beginning of the experiment, water is passed through the refrigerator and the flask is slightly heated, carefully observing the temperature and pressure at which the water begins to boil. A very strong vacuum in this experiment should not be allowed in order to avoid cracking of the tubes.
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 A simplified version of the experiment: we heat the water in the flask to a boil, remove the flask from the stove and close it hermetically with a stopper - boiling stops, we place the flask under a stream of cold water - rapid boiling resumes. |
An experience. Boiling water at a pressure above atmospheric pressure. The device is assembled in accordance with.
The flask for the device is taken wide-mouthed, round-bottomed, made of thick and high-quality glass with a capacity of 500 ml.
Pour 250 ml of pre-boiled water into the flask. The flask is fixed in a tripod and closed with a rubber stopper through which two glass tubes are passed. One tube, 6-7 mm in diameter, ends with a bubble of such a size that it passes through the neck of the flask. The second tube, 6 mm in diameter, starts at the bottom edge of the cork; outside, it is bent at an angle of 90 ° and, using a thick-walled rubber tube, is connected to another glass tube bent at a right angle, lowered almost to the bottom into a cylinder with mercury 90-100 cm high and 1.5-2 cm in diameter.
Several pieces of porous porcelain are placed in a vial and filled up to half with water.
With the indicated amount of mercury, the air in the flask is under a pressure of more than two atmospheres.
So that the tube lowered into the cylinder with mercury is not thrown out, it is fixed in the tripod clamp.
After assembling the device, heat the flask with water. At the beginning, water in a bubble boils under atmospheric pressure, and much later water in a flask boils under a pressure of more than two atmospheres.
Round-bottom flasks are used for the experiment, as they are more resistant to high pressure.
During the experiment, they work carefully, observing at a certain distance, since at a pressure of 2-3 atm the flask may burst.
Water is involved in the following chemical reactions: in reactions in which it exhibits oxidizing properties, in reactions of hydrolysis, hydration, addition, substitution, and in reactions in which water plays the role of a catalyst.
In experiments on the production of hydrogen, the oxidative effect of water on sodium, potassium, calcium, magnesium, aluminum, iron and carbon was considered.
The sections devoted to bromine and iodine describe experiments on the production of hydrogen bromide and iodide by hydrolysis of phosphorus halides.
When considering the properties of chlorine, bromine and hydrogen chloride, hydration was discussed, which proceeds as an addition reaction.
In experiments illustrating the combination of hydrogen with chlorine or iodine with zinc, the catalytic properties of water are shown.
Chemical reactions involving water occur in many of the experiments described.
Substances enter into chemical reactions, as a result of which other substances are formed. Are there any changes in the mass of the substance as a result of the reaction? Scholars have made various theories on this issue.
The famous English chemist R. Boyle, calcining various metals in an open retort and weighing them before and after heating, found that the mass of metals becomes larger. Based on these experiments, he did not take into account the role of air and made the wrong conclusion that the mass of substances as a result chemical reactions changes. R. Boyle argued that there is some kind of “fiery matter”, which, in the case of heating the metal, combines with the metal, increasing the mass.
M. V. Lomonosov, unlike R. Boyle, calcined metals not in the open air, but in sealed retorts and weighed them before and after calcination. (A retort with a brazier is shown in Fig. 35, see p. 54.) He proved that the mass of substances before and after the reaction remains unchanged and that when calcined, some air is added to the metal. (Oxygen had not yet been discovered at that time.) He formulated the results of these experiments in the form of a law: “All changes that occur in nature are such a state of affairs that how much is taken from one body, so much will be added to another.” This law is currently formulated as follows:
The mass of substances that entered into a chemical reaction is equal to the mass of the formed substances.
Much later (1789), the law of conservation of mass was independently established by M. V. Lomonosov by the French chemist A. Lavoisier (p. 55).
It is possible to confirm the correctness of the law of conservation of mass of substances by a simple experiment. A little red phosphorus is placed in the flask (Fig. 16), closed with a cork and weighed on a balance (a). Then the flask with phosphorus (b) is gently heated. The fact that a chemical reaction has occurred is judged by the appearance of white smoke in the flask, consisting of particles of phosphorus (V) oxide. During the second weighing, they are convinced that as a result of the reaction the mass of substances has not changed (c).
From the point of view of atomic and molecular theory, the law of conservation of mass is explained as follows: as a result of chemical reactions, atoms do not disappear and do not appear, but their rearrangement occurs.
Since the number of atoms before and after the reaction remains unchanged, their total mass also does not change.
The meaning of the law of conservation of mass of substances.
1. The discovery of the law of conservation of mass of substances contributed to further development chemistry as a science.
2. On the basis of the law of conservation of mass of substances, practically important calculations are made. For example, you can calculate how many starting materials are required to obtain iron (II) sulfide weighing 44 kg if iron and sulfur react in a mass ratio of 7:4. According to the law of conservation of mass of substances, the interaction of iron weighing 7 kg and sulfur weighing 4 kg produces iron (II) sulfide weighing 11 kg. And since it is necessary to obtain iron (II) sulfide weighing 44 kg, i.e. 4 times more, then 4 times more starting materials will also be required: 28 kg of iron (7-4) and 16 kg of sulfur (4-4 ).
3. On the basis of the law of conservation of mass of substances, the equations of chemical reactions are compiled.
Answer questions 1-3 (page 42).
§15. Chemical Equations
A chemical equation is a conditional record of a chemical reaction by means of chemical signs and formulas.
According to the chemical equation of reactions, one can judge which substances react and which are formed. When compiling reaction equations, proceed as follows:
1. On the left side of the equation, write the formulas of the substances that react, and then put an arrow. It must be remembered that the molecules of simple gaseous substances almost always consist of two atoms (O 2, H 2, C1 2, etc.):
2. On the right side (after the arrow) write the formulas of the substances formed as a result of the reaction:
3. The reaction equation is compiled on the basis of the law of conservation of mass of substances, i.e., on the left and on the right there should be the same number of atoms. This is achieved by placing the coefficients in front of the formulas of substances. First, the number of atoms is equalized, of which there are more in the reacting substances. In our examples, these are oxygen atoms. Find the least common multiple of the number of oxygen atoms in the left and right parts of the record from the arrow. In the reaction of magnesium with oxygen, the least common multiple is the number 2, and in the example with phosphorus, the number is 10. When dividing the least common multiple by the number of corresponding atoms (in the examples given, by the number of oxygen atoms), on the left and right sides of the record from the arrow, one finds corresponding coefficients as shown in the following diagram:
Equalize the number of atoms of other chemical elements. In our examples, the number of magnesium and phosphorus atoms should be equalized:
In those cases when, when compiling chemical equations, the thermal effects of reactions are not indicated, an arrow is placed instead of the equal sign.
§ 16. Types of chemical reactions
chemical reactions
can be classified into four main types: 1)
decomposition;
2)
connections;
3)
substitution;
4)
exchange
(p. 82).
You got acquainted with the decomposition reaction using the example of water decomposition (p. 13). You know the compound reaction from the example of the interaction of sulfur with iron (p. 15).
To get acquainted with the substitution reaction, you can perform the following experiment. A cleaned iron nail (or iron filings) is lowered into a blue solution of copper (II) chloride CuCl 2 . The nail (sawdust) is immediately covered with a coating of copper, and the solution turns from blue to greenish, since instead of copper (II) chloride СuС1 2, iron chloride (II) FeCl 2 is formed. The ongoing chemical reaction is expressed by the chemical equation
Fe + CuCl 2 ->Cu + FeCl 2
When comparing the chemical reactions discussed above, one can give them definitions and reveal their features (Scheme 6).
1 You will get acquainted with exchange reactions in a further course in chemistry (p. 82).
2 In order to start the reaction, in many cases heating is required. Then, in the reaction equations, the sign t is placed above the arrow.
3 If a gas is released as a result of the reaction, an arrow is placed next to its formula Beepx, and if a substance precipitates, then an arrow down is placed next to the formula of this substance.
Do exercises 5-7 (pp. 42-43).
1. By whom, when and how was the law of conservation of mass discovered? Give the formulation of the law and explain it from the point of view of atomic and molecular theory.
2. Zinc powder was poured into the retort (Fig. 35), the gas outlet tube was closed with a clamp, the retort was weighed, and the contents were calcined. When the retort had cooled down, it was weighed again. Has its mass changed and why? Then the clamp was opened. Have the scales remained in balance and why?
3. What theoretical and practical value has the law of conservation of mass of matter? Give examples.
4. Adhering to the sequence previously given (see p. 35), and taking into account the valency of the elements, compose the reaction equations according to the following schemes:
5. Write two reaction equations for each of the types you know and explain their essence from the point of view of atomic and molecular theory.
6. Given metals: calcium Ca, aluminumAI, lithiumLi. Make up the equations of chemical reactions of these metals with oxygen, chlorine and sulfur, if it is known that sulfur in compounds with metals and hydrogen is divalent.
7. Rewrite the schemes of reaction equations below, instead of question marks, write the formulas of the corresponding substances, arrange the coefficients and explain what type each of the indicated reactions belongs to:
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