A salt can be defined as a compound that is formed by the reaction between an acid and a base, but is not water. This section will consider those properties of salts that are associated with ionic equilibria.
reactions of salts in water
It will be shown a little later that solubility is a relative concept. However, for the purposes of the discussion ahead, we can roughly divide all salts into those that are soluble and those that are insoluble in water.
Some salts form neutral solutions when dissolved in water. Other salts form acidic or alkaline solutions. This is due to the occurrence of a reversible reaction between salt ions and water, as a result of which conjugate acids or bases are formed. Whether the salt solution turns out to be neutral, acidic or alkaline depends on the type of salt. In this sense, there are four types of salts.
Salts formed by strong acids and weak bases. Salts of this type, when dissolved in water, form an acidic solution. Let us take ammonium chloride NH4Cl as an example. When this salt is dissolved in water, the ammonium ion acts as
The excess amount of H3O+ ions formed in this process causes the acidic properties of the solution.
Salts formed by a weak acid and a strong base. Salts of this type, when dissolved in water, form an alkaline solution. As an example, let's take sodium acetate CH3COONa1. The acetate ion acts as a base, accepting a proton from water, which in this case acts as an acid:
The excess amount of OH- ions formed in this process determines the alkaline properties of the solution.
Salts formed by strong acids and strong bases. When salts of this type are dissolved in water, a neutral solution is formed. Let's take sodium chloride NaCl as an example. When dissolved in water, this salt is completely ionized, and, therefore, the concentration of Na+ ions turns out to be equal to the concentration of Cl- ions. Since neither one nor the other ion enters into acid-base reactions with water, an excess amount of H3O+ or OH ions does not form in the solution. Therefore, the solution turns out to be neutral.
Salts formed by weak acids and weak bases. An example of this type of salt is ammonium acetate. When dissolved in water, ammonium ion reacts with water as an acid, and acetate ion reacts with water as a base. Both of these reactions are described above. An aqueous solution of a salt formed by a weak acid and a weak base can be weakly acidic, weakly alkaline, or neutral, depending on the relative concentrations of the H3O+ and OH- ions formed as a result of the reactions of the salt's cations and anions with water. This depends on the relationship between the values of the dissociation constants of the cation and anion.
Common salt - sodium chloride - is perhaps the most valuable food product. And not only because we cannot live without the elements that it consists of - sodium and chlorine - but also because salty taste is one of the main taste sensations. Salt not only has its own flavor, but it can also magically enhance or enhance other taste sensations.
The word "salt" does not mean any one substance. In chemical terms, it is a general designation for an entire family of chemicals. In terminology, a salt is a reaction product between an acid and an alkali.
Some other types of salt used in gastronomy include potassium chloride, which serves as a salt substitute in low-salt diets; potassium iodide, which is added to salt to provide iodine in our diet; and finally, sodium nitrite - used along with sodium nitrate - necessary for salting various meat products.
If there are so many different types of salts, can we say that salinity is a unique characteristic of sodium chloride? This is wrong. Try one of the potassium chloride "salt substitutes" and you will describe it as "salty," but that saltiness is not the same as that familiar taste of sodium chloride - just as the sensation of sweetness is slightly different with different types of sugars and artificial ones. sweeteners.
Salt has been used for thousands of years not only as a nutrient and seasoning, but also as a preservative for meat, fish and vegetables, which, thanks to salting, could be eaten not immediately after the end of the hunt or harvest, but much later.
What types of salt are there?
The number of types of specialty salt is simply staggering. Manufacturers today produce about 60 types for the food industry and the average consumer, including flake and fine-flake salt, coarse, fine, ultra-fine and finely ground salt. Chemically, they all contain more than 99% sodium chloride, but have different physical characteristics for use in the preparation of various foods - from potato chips, popcorn, toasted nuts to pies, various types of bread, cheese, crackers, margarine, peanuts butter and pickles.
For a margarita, you'll want large crystals that will stick to the lime juice on the rim of the glass, because smaller salt crystals will simply dissolve in the juice. For popcorn, on the other hand, you want the exact opposite: flour-like crystals that will get into the cracks of the corn kernels and stay there.
What is the difference between sea and “regular” salt
When we hear names like sea salt and regular salt (or table salt), we might think that they refer to two different substances with different properties. But it's not that simple. Salt actually comes from two different sources: underground mines and sea water. But this fact alone does not make them fundamentally different.
We inherited underground salt deposits from dried up ancient seas that disappeared at one stage or another in the history of our planet - from several million to hundreds of millions of years ago. Then, thanks to geological processes, some salt deposits were closer to the surface of the earth, and now they exist in the form of peculiar “domes”. Other salt deposits are hundreds of meters deeper and therefore more difficult to mine.
Rock salt is crushed by large machines in cavities cut out of the salt masses. But rock salt is not suitable for human consumption because when the ancient seas dried out, they retained silt and various organic remains. Therefore, table salt is mined differently: water is pumped into the mine shaft to dissolve the salt, salt water (saline solution) is pumped to the surface, all impurities are settled and, finally, the now pure salt solution is evaporated using a vacuum. The result is the familiar tiny crystals of table salt.
In coastal areas where sunny weather prevails, salt can be obtained by allowing the sun and wind to evaporate water from shallow ponds or “islands.”
Is sea salt good for you?
If you evaporate all the water from the ocean (after removing the fish from there), you will be left with a sticky, gray and bitter mass of silt, 78% consisting of sodium chloride - ordinary salt. The remaining 22% consists of 99% magnesium and calcium compounds, which are responsible for bitterness. In addition, at least 75 other chemical elements are present in very small quantities. This last fact is the basis for widespread statements about the “mass of nutrient minerals” in sea salt.
However, chemical analysis will dampen our enthusiasm: minerals, even in such raw and untreated sludge, are present in insignificant quantities. For example, you would have to eat two tablespoons of this to get the amount of iron you get from a single grape.
The idea that sea salt already contains iodine is a myth. Because certain types of marine vegetation are rich in iodine, some people consider the ocean to be a kind of “iodized soup.” As for the chemical elements present in seawater, it contains 100 times more boron than iodine, but I have never heard sea salt advertised as a source of boron.
What is in store-bought salt, besides the salt itself?
It is often written about sea salt that it does not contain “additives with an unpleasant taste”, like table salt. However, whatever its origin, salt in any case contains anti-caking additives (for example, calcium silicate) so that its granules fall off easily; Salt crystals are essentially little cubes, and they tend to stick to each other. Due to the fact that calcium silicate (like all other anti-caking additives) does not dissolve in water, table salt, when dissolved in water, produces a whitish solution.
Other anti-caking agents include magnesium carbonate (E504), calcium carbonate (E170) and calcium phosphates (E341). All these chemicals are tasteless and odorless. But even if they had taste and smell and professional tasters could distinguish the subtlest shades of taste in solid salt that arose due to the introduction of these additives (in a volume of less than 1%), there would still be a dilution factor that occurs when adding salt according to any recipe , would not allow the tasters to achieve their goal.
Does salt taste different?
Depending on how the salt was collected and processed, the crystals of different brands of sea salt can vary greatly in shape, from flakes to pyramids to irregularly shaped fragments (you can see this if you take a magnifying glass). The size of the crystals also varies, from very small to large, although all are larger than regular table salt.
If this salt is sprinkled on relatively dry food, such as a slice of tomato, the larger, flakier crystals can create small patches of saltiness - when they touch the tongue and then dissolve, or when they land on the teeth and are crushed. That's why chefs value sea salt so much: for those little "flashes" of salty flavor. Table salt is incapable of this, since its compact small crystals dissolve on the tongue much more slowly. Thus, it is the complex shape of the crystals, and not their marine origin, that determines the taste characteristics of many types of sea salt.
“What Einstein told his cook. Physics and chemistry in your kitchen" Robert Wolke
Food is something ordinary for us, we rarely think about what and how we eat, what happens to dishes and products before they appear on our table, why we like some of them more, others less, why some Some of them are useful, while others are not.
Meanwhile, every day miracles happen in the kitchen that we do not notice. The author clearly, simply and wittily explains their nature and background. The main content of the book consists of everyday questions, to which the author provides answers, explaining them from a scientific point of view, in a popular, accessible form.
In this book, the author answers more than a hundred questions that readers of his Washington Post column, including home-cookers and professional chefs, have asked him over the years: Why is sugar sweet? Why does chocolate melt in your mouth? How is coffee decaffeinated? How much alcohol is in alcoholic drinks? And much more.
Maria Rodenko
Experiments with water for preschoolers
Sorceress water
Water, steam, ice are the same substance!
Pour some water into a saucer and put it in the freezer for 2-3 hours. What happened?
Place a saucer of ice on the table. How long will it take for it to contain water again? What happens to ice - solid water?
Leave a saucer of water on the windowsill for 2-3 days. Will it evaporate soon?
Explanation: water, steam, ice are the same substance, but only in different states: liquid, solid and gaseous.
2. Is it possible to glue paper with water?
Take two sheets of paper, attach them to each other and try to move them in different directions. Happened?
Then wet the sheets of paper with water, attach them to each other, press lightly and try to move them, one in one direction, the other in the other.
Explanation: Water has a “gluing” effect, so paper can be glued together with water.
3. Is it possible to speed up the evaporation of water?
Wet your hands with tap water and wave them quickly. How do your hands feel? Why is this happening?
Explanation: The evaporation of water can be accelerated, for example, by creating air movement. At the same time, water particles evaporate faster, taking with them more heat. Therefore, your hands feel cool when you wave your arms.
Droplets - balls
Take fine sand or flour and sprinkle with water. Droplets are formed - balls.
Spray the leaves of the plant with a spray bottle. What kind of droplets did you get?
Explanation: the particles collect droplets of water around themselves and form one large droplet-ball, and when many droplets of water fall on the leaf of a plant, they gather together to form a large droplet-ball or even a small puddle.
5. Sugar dissolves in water.
Place a piece of sugar in a glass and pour the tea in a thin stream, trying to get only the sugar. The sugar will gradually erode and then... disappear? Of course not.
Use a spoon to scoop a little tea from a glass and a spoonful of tea from a teapot, taste and compare the taste. What can you smell, does the tea taste the same?
Explanation: Sugar dissolves in water and mixes with it, so the water becomes sweet.
6. Salt dissolves in water.
Pour one tablespoon of salt into a glass of water and stir.
What happened? Has the salt “disappeared”? Let the child taste some water. What did the water become?
Explanation: the salt did not disappear, it dissolved in the water, and the water became salty.
The salt evaporates and crystallizes.
Pour 2-3 tablespoons of salt into a glass of water. Stir until completely dissolved. Then place it in a sunny place and observe the behavior of the salt.
After a few days, salt crystals will appear on the walls of the glass as the water evaporates.
Explanation: the water evaporates and salt crystals settle on the walls of the glass.
8. Sand does not dissolve.
Invite your child to compare sugar and sand, find out what dissolves in water and what does not.
Stir a spoonful of river sand in one glass of water and a spoonful of sugar in another. Let it sit.
See what happened to grains of sand and water.
Explanation: the water with river sand has become cloudy and dirty. Heavier grains of sand sank to the bottom, and smaller ones float in the water, making it cloudy. Sugar became invisible in water, water dissolved it.
9. Pipette tube.
Take two identical glasses, one with water and the other empty. Place the cocktail straw in a glass of water, hold it on top with your index finger and, without lifting your finger, transfer it to an empty glass. Remove your finger from the straw and water will flow out.
After doing the same several times, you can transfer all the water from one glass to another.
Swimming fish.
Draw and cut out a fish on cardboard. Pour water into a basin. Dip a toothpick into liquid soap or vegetable oil and place a dot on the tail of the fish.
Place the fish on the water so that the tail is at the edge of the pelvis, as a result the fish will swim.
To repeat the experiment, you need to wash and dry the basin.
Explanation: a drop of oil spreads over the water, its particles are repelled by water particles, and as a result the fish swims.
11. Floating egg.
Fill a liter jar halfway with water. Using a spoon, carefully lower the raw egg into it and remove the spoon. How will the egg behave?
Carefully remove the egg. Pour half a cup (125 ml) of salt into the same jar of water and stir with a spoon. Then place the egg in water. How does the egg behave now?
Explanation: An egg sinks in clear water because it is heavier than water. By adding salt to water, we increase its density, and therefore the egg floats in it.
12. Singing bottles. High and low sounds.
Fill 3 identical bottles halfway with water, and then pour half of the water from one to the other. Bring the bottle to your lips and blow over the neck to hear the singing sound. Blow over other bottles, do they sing the same?
Arrange the bottles in order of pitch.
Explanation: The exhaled air above the bottle causes the air inside it to vibrate. The pitch of the sound depends on the amount of air inside the bottle.
13. Rainbow paper.
Fill a deep bowl halfway with water. Gently add a drop of clear nail polish; the polish will spread across the surface of the water.
Immerse a piece of black double-sided paper in water, remove it and dry it. You can see rainbow stains on dry paper.
Explanation: The varnish forms a thin film on the surface of the water. The film envelops the paper, and the light hitting it is reflected, forming a rainbow pattern.
14. Blooming flowers.
Draw and cut out flowers with long petals from colored paper. Using a pencil, curl the petals towards the center of the flower.
Pour water into a basin and place the flowers on it. They will begin to straighten their petals and bloom.
Explanation: when in contact with water, the paper gets wet, becomes heavier, and the flower's petals begin to bloom.
15. Water not spilling out.
Pour a glass of water to the brim. Place a postcard or square of thick cardboard on top. Press the card onto the glass with your hand and turn it upside down over the sink.
Remove your hand. What happened?
Explanation: the card does not fall and the water does not pour out because the air in the glass presses on it from below and presses it to the edge of the glass. Water will spill out if the card is moved.
16. Invisible ink.
Squeeze the juice from half a lemon into a cup and add the same amount of water. Dip a cotton swab into the prepared solution. Write something for her on a piece of paper.
When the inscription is dry, heat the paper over the switched on table lamp. Previously invisible words will appear on paper.
17.
Jumping grains.
Pour sparkling water into a glass and add 6 grains of rice.
Wait a few seconds and watch through the glass as the grains begin to jump.
Explanation: Rice is slightly heavier than water, when it hits the glass it begins to sink. Gas bubbles land on it and rise up. The bubbles on the surface burst and the gas evaporates. The heavy rice falls down again. It will go down and up until the water runs out.
Back forward
Attention! Slide previews are for informational purposes only and may not represent all the features of the presentation. If you are interested in this work, please download the full version.
The purpose of the lesson: study of the properties of water.
Lesson objectives: give an idea of water as a solvent, of soluble and insoluble substances; introduce the concept of “filter”, the simplest methods for determining soluble and insoluble substances; prepare a report on the topic “Water is a solvent.”
Equipment and visual aids: textbooks, reading books, notebooks for independent work; sets: glasses empty and with boiled water; boxes with table salt, sugar, river sand, clay; teaspoons, funnels, paper napkin filters; gouache (watercolor paints), brushes and reflection sheets; presentation made in Power Point, multimedia projector, screen.
DURING THE CLASSES
I. Organizational moment
U. Good morning everyone! (Slide 1)
I invite you to the third meeting of the school science club “We and the world around us.”
II. Communicating the topic and purpose of the lesson
Teacher. Today we have guests, teachers from other schools who came to the club meeting. I propose to the chairman of the club, Anastasia Poroshina, to open the meeting.
Chairman. Today we gathered for a club meeting on the topic “Water is a solvent.” The task for all those present is to prepare a report on the topic “Water is a solvent.” In this lesson you will once again become researchers of the properties of water. You will study these properties in your laboratories, with the help of “consultants” – Mikhail Makarenkov, Olesya Starkova and Yulia Stenina. Each laboratory will have to complete the following task: conduct experiments and observations, and at the end of the meeting discuss the plan for the “Water - Solvent” message.
III. Learning new material
U. With the chairman's permission, I would like to make my first announcement. (Slide 2) The same meeting on the topic “Water is a solvent” was recently held by students from the village of Mirny. The meeting was opened by Kostya Pogodin, who reminded everyone present of another amazing property of water: many substances in water can disintegrate into invisible tiny particles, that is, dissolve. Therefore, water is a good solvent for many substances. After this, Masha proposed to conduct experiments and identify methods by which it would be possible to obtain an answer to the question of whether a substance dissolves in water or not.
U. At a club meeting, I suggest that you determine the solubility in water of substances such as table salt, sugar, river sand and clay.
Let's assume which substance, in your opinion, will dissolve in water and which will not dissolve. Express your assumptions, guesses and continue your statement: (Slide 3)
U. Let's think together about what hypotheses we will confirm. (Slide 3)
Suppose... (salt dissolves in water)
Let's say... (sugar will dissolve in water)
Perhaps... (sand will not dissolve in water)
What if... (clay will not dissolve in water)
U. Come on, let's conduct experiments that will help us figure this out. Before work, the chairman will remind you of the rules for conducting experiments and hand out cards on which these rules are printed. (Slide 4)
P. Look at the screen where the rules are written.
"Rules for conducting experiments"
- All devices must be handled with care. Not only can they be broken, they can also cause injury.
- While working, you can not only sit, but also stand.
- The experiment is carried out by one of the students (the speaker), the rest silently observe or, at the request of the speaker, help him.
- The exchange of opinions on the results of the experiment begins only after the speaker allows it to begin.
- You need to talk to each other quietly, without disturbing others.
- Approaching the table and changing laboratory equipment is only possible with the permission of the chairman.
IV. Practical work
U. I suggest that the chairman choose a “consultant” who will read aloud from the textbook (p. 85) the procedure for conducting the first experiment. (Slide 5)
1)
P. Swipe experiment with table salt. Check if table salt dissolves in water.
A “consultant” from each laboratory takes one of the prepared sets and conducts an experiment with table salt. Boiled water is poured into a transparent glass. Pour a small amount of table salt into the water. The group observes what happens to the salt crystals and tastes the water.
The chairman (as in the KVN game) reads the same question to each group, and representatives from the laboratories answer them.
P.(Slide 6) Has the clarity of the water changed? (Transparency has not changed)
Has the color of the water changed? (Color has not changed)
Has the taste of the water changed? (The water has become salty)
Can we say that the salt has disappeared? (Yes, she dissolved, disappeared, she is not visible)
U. Draw a conclusion. (Salt has dissolved)(Slide 6)
P. I ask everyone to proceed with the second experiment, for which it is necessary to use filters.
U. What is a filter? (A device, device or structure for purifying liquids, gases from solid particles and impurities.)(Slide 7)
U. Read aloud the procedure for performing the filter experiment. (Slide 8)
Students pass water with salt through a filter, observe and taste the water.
P.(Slide 9) Is there any salt left on the filter? (No table salt remains on the filter)
Has the taste of the water changed? (The taste of the water has not changed)
Have you managed to clear the salt from the water? (Table salt passed through the filter with water)
U. Draw a conclusion from your observations. (Salt dissolved in water)(Slide 9)
U. Was your hypothesis confirmed?
U. Everything is correct! Well done!
U. Prepare the results of the experiment in writing in your Notebook for independent work (p. 30). (Slide 10)
2)
P.(Slide 11) Let's do the same experience again, but instead of salt we put a teaspoon granulated sugar.
The “consultant” from each laboratory takes the second set and conducts an experiment with sugar. Boiled water is poured into a transparent glass. Add a small amount of sugar to the water. The group observes what happens and tastes the water.
P.(Slide 12) Has the transparency of the water changed? (Water clarity has not changed)
Has the color of the water changed? (The color of the water has not changed)
Has the taste of the water changed? (The water has become sweet)
Can we say that sugar has disappeared? (Sugar became invisible in water, water dissolved it)
U. Draw a conclusion. (Sugar has dissolved)(Slide 12)
U. Pass the water and sugar through a paper filter. (Slide 13)
Students pass water with sugar through a filter, observe and taste the water.
P.(Slide 14) Is there any sugar left on the filter? (Sugar is not visible on the filter)
Has the taste of the water changed? (The taste of the water has not changed)
Have you managed to remove sugar from water? (It was not possible to purify the water from sugar; it went through the filter along with the water)
U. Draw a conclusion. (Sugar dissolved in water)(Slide 14)
U. Was the hypothesis confirmed?
U. Right. Well done!
U. Prepare the results of the experiment in writing in your Notebook for independent work. (Slide 15)
3) P.(Slide 16) Let's check the statements and conduct river sand experience.
U. Read the procedure for conducting the experiment in the textbook.
Conduct an experiment with river sand. Stir a teaspoon of river sand in a glass of water. Let the mixture settle. Observe what happens to grains of sand and water.
P.(Slide 17) Has the transparency of the water changed? (The water has become cloudy and dirty)
Has the color of the water changed? (The color of the water has changed)
Have the grains of sand disappeared? (Heavier grains of sand sink to the bottom, and smaller ones float in the water, making it cloudy)
U. Draw a conclusion. (The sand did not dissolve)(Slide 17)
U.(Slide 18) Pass the contents of the glass through a paper filter.
Students pass water with sugar through a filter and observe.
P.(Slide 19) What passes through the filter and what remains on it? (The water passes through the filter, but the river sand remains on the filter and the grains of sand are clearly visible)
Has the water been cleared of sand? (The filter helps clean the water from particles that do not dissolve in it)
U. Draw a conclusion. (River sand did not dissolve in water)(Slide 19)
U. Was your assumption about the solubility of sand in water correct?
U. Great! Well done!
U. Prepare the results of the experiment in writing in your Notebook for independent work. (Slide 20)
4) P.(Slide 21) Do the same experiment with a piece of clay.
Conduct an experiment with clay. Stir a piece of clay in a glass of water. Let the mixture settle. Observe what happens to clay and water.
P.(Slide 22) Has the transparency of the water changed? (The water has become cloudy)
Has the color of the water changed? (Yes)
Have the clay particles disappeared? (Heavier particles sink to the bottom, and smaller ones float in the water, making it cloudy)
U. Draw a conclusion. (The clay did not dissolve in water)(Slide 22)
U.(Slide 23) Pass the contents of the glass through a paper filter.
P.(Slide 24) What passes through the filter, and what remains on it? (Water passes through the filter, and undissolved particles remain on the filter.)
Has the water been cleared of clay? (The filter helped clear the water of particles that did not dissolve in the water)
U. Draw a conclusion. (Clay does not dissolve in water)(Slide 24)
U. Was the hypothesis confirmed?
U. Well done! Everything is correct!
U. I ask one of the group members to read the conclusions written in the notebook to everyone present.
U. Does anyone have any additions or clarifications?
U. Let's draw conclusions from the experiments. (Slide 25)
Are all substances soluble in water? (Salt and granulated sugar dissolved in water, but sand and clay did not dissolve.)
Is it always possible to use a filter to determine whether a substance dissolves in water or not? (Substances dissolved in water pass through the filter along with the water, and undissolved particles remain on the filter)
U. Read about the solubility of substances in water in the textbook (p. 87).
U. Draw a conclusion about the properties of water as a solvent. (Water is a solvent, but not all substances dissolve in it)(Slide 25)
U. I advise club members to read the story in the anthology “Water is a Solvent” (p. 46). (Slide 26)
Why have scientists not yet been able to obtain absolutely pure water? (Because there are hundreds, and maybe thousands of different substances dissolved in water)
U. How do people use the ability of water to dissolve certain substances?
(Slide 27) Tasteless water becomes sweet or salty thanks to sugar or salt, as water dissolves and acquires their taste. A person uses this property when preparing food: brewing tea, making compote, soups, salting and canning vegetables, making jam.
(Slide 28) When we wash our hands, wash or bathe, when we wash clothes, we use liquid water and its properties as a solvent.
(Slide 29) Gases, in particular oxygen, also dissolve in water. Thanks to this, fish and others live in rivers, lakes, and seas. In contact with air, water dissolves oxygen, carbon dioxide and other gases that are in it. For living organisms living in water, such as fish, oxygen dissolved in water is very important. They need it to breathe. If oxygen did not dissolve in water, then bodies of water would be lifeless. Knowing this, people do not forget to saturate the water in the aquarium where the fish live with oxygen, or cut ice holes in reservoirs in winter to improve life under the ice.
(Slide 30) When we paint with watercolors or gouache.
U. Pay attention to the task written on the board. (Slide 31) I propose to draw up a collective plan for a presentation on the topic “Water is a solvent.” Discuss it in your laboratories.
Listening to plans on the topic “Water is a solvent” compiled by students.
U. Let's all formulate a plan for the speech together. (Slide 31)
Sample plan for a speech on the topic “Water is a solvent”
- Introduction.
- Dissolution of substances in water.
- Conclusions.
- People use the properties of water to dissolve certain substances.
Excursion to the Exhibition Hall.(Slide 32)
U. When preparing your message, you can use additional literature selected by the guys, assistant speakers on the topic of our meeting. (Draw students’ attention to the exhibition of books and Internet pages)
V. Lesson summary
What property of water was studied at the club meeting? (Property of water as a solvent)
What conclusion did we come to after studying this property of water? (Water is a good solvent for some substances.)
Do you think it is difficult to be researchers?
What did you find most challenging or interesting?
Will the knowledge acquired during the study of this property of water be useful to you in later life? (Slide 33) (It is very important to remember that water is a solvent. Water dissolves salts, some of which are both beneficial for humans and harmful. Therefore, you cannot drink water from a source if you do not know whether it is pure. It is not for nothing that people eat proverb: “Not all water is suitable for drinking.”)
VI. Reflection
How do we use the ability of water to dissolve certain substances in art lessons? (When we paint with watercolors or gouache)
I suggest you, using this property of water, paint the water in a glass in a color that best suits your mood. (Slide 34)
“Yellow color” – joyful, bright, good mood.
“Green color” – calm, balanced.
“Blue color” is a sad, melancholy, melancholy mood.
Display your sheets with colored water in a glass.
VII. Assessment
I thank the chairman, “consultants” and all participants of the meeting for their active work.
VIII. Homework
| Cations | Anions | |||||||||
| F- | Cl- | Br- | I - | S 2- | NO 3 - | CO 3 2- | SiO 3 2- | SO 4 2- | PO 4 3- | |
| Na+ | R | R | R | R | R | R | R | R | R | R |
| K+ | R | R | R | R | R | R | R | R | R | R |
| NH4+ | R | R | R | R | R | R | R | R | R | R |
| Mg 2+ | RK | R | R | R | M | R | N | RK | R | RK |
| Ca2+ | NK | R | R | R | M | R | N | RK | M | RK |
| Sr 2+ | NK | R | R | R | R | R | N | RK | RK | RK |
| Ba 2+ | RK | R | R | R | R | R | N | RK | NK | RK |
| Sn 2+ | R | R | R | M | RK | R | N | N | R | N |
| Pb 2+ | N | M | M | M | RK | R | N | N | N | N |
| Al 3+ | M | R | R | R | G | R | G | NK | R | RK |
| Cr 3+ | R | R | R | R | G | R | G | N | R | RK |
| Mn 2+ | R | R | R | R | N | R | N | N | R | N |
| Fe 2+ | M | R | R | R | N | R | N | N | R | N |
| Fe 3+ | R | R | R | - | - | R | G | N | R | RK |
| Co2+ | M | R | R | R | N | R | N | N | R | N |
| Ni 2+ | M | R | R | R | RK | R | N | N | R | N |
| Cu 2+ | M | R | R | - | N | R | G | N | R | N |
| Zn 2+ | M | R | R | R | RK | R | N | N | R | N |
| Cd 2+ | R | R | R | R | RK | R | N | N | R | N |
| Hg 2+ | R | R | M | NK | NK | R | N | N | R | N |
| Hg 2 2+ | R | NK | NK | NK | RK | R | N | N | M | N |
| Ag+ | R | NK | NK | NK | NK | R | N | N | M | N |
Legend:
P - the substance is highly soluble in water; M - slightly soluble; H - practically insoluble in water, but easily soluble in weak or dilute acids; RK - insoluble in water and soluble only in strong inorganic acids; NK - insoluble in either water or acids; G - completely hydrolyzes when dissolved and does not exist in contact with water. A dash means that such a substance does not exist at all.
In aqueous solutions, salts completely or partially dissociate into ions. Salts of weak acids and/or weak bases undergo hydrolysis. Aqueous solutions of salts contain hydrated ions, ion pairs and more complex chemical forms, including hydrolysis products, etc. A number of salts are also soluble in alcohols, acetone, acid amides and other organic solvents.
From aqueous solutions, salts can crystallize in the form of crystal hydrates, from non-aqueous solutions - in the form of crystal solvates, for example CaBr 2 3C 2 H 5 OH.
Data on various processes occurring in water-salt systems, on the solubility of salts in their joint presence depending on temperature, pressure and concentration, on the composition of solid and liquid phases can be obtained by studying the solubility diagrams of water-salt systems.
General methods for the synthesis of salts.
1. Obtaining medium salts:
1) metal with non-metal: 2Na + Cl 2 = 2NaCl
2) metal with acid: Zn + 2HCl = ZnCl 2 + H 2
3) metal with a salt solution of a less active metal Fe + CuSO 4 = FeSO 4 + Cu
4) basic oxide with acidic oxide: MgO + CO 2 = MgCO 3
5) basic oxide with acid CuO + H 2 SO 4 = CuSO 4 + H 2 O
6) bases with acid oxide Ba(OH) 2 + CO 2 = BaCO 3 + H 2 O
7) bases with acid: Ca(OH) 2 + 2HCl = CaCl 2 + 2H 2 O
8) salts with acid: MgCO 3 + 2HCl = MgCl 2 + H 2 O + CO 2
BaCl 2 + H 2 SO 4 = BaSO 4 + 2HCl
9) base solution with salt solution: Ba(OH) 2 + Na 2 SO 4 = 2NaOH + BaSO 4
10) solutions of two salts 3CaCl 2 + 2Na 3 PO 4 = Ca 3 (PO 4) 2 + 6NaCl
2.Obtaining acid salts:
1. Interaction of an acid with a lack of base. KOH + H2SO4 = KHSO4 + H2O
2. Interaction of the base with excess acid oxide
Ca(OH) 2 + 2CO 2 = Ca(HCO 3) 2
3. Interaction of the average salt with the acid Ca 3 (PO 4) 2 + 4H 3 PO 4 = 3Ca(H 2 PO 4) 2
3.Obtaining basic salts:
1. Hydrolysis of salts formed by a weak base and a strong acid
ZnCl 2 + H 2 O = Cl + HCl
2. Adding (drop by drop) small amounts of alkalis to solutions of medium metal salts AlCl 3 + 2NaOH = Cl + 2NaCl
3. Interaction of salts of weak acids with medium salts
2MgCl 2 + 2Na 2 CO 3 + H 2 O = 2 CO 3 + CO 2 + 4NaCl
4. Obtaining complex salts:
1. Reactions of salts with ligands: AgCl + 2NH 3 = Cl
FeCl 3 + 6KCN] = K 3 + 3KCl
5. Obtaining double salts:
1. Joint crystallization of two salts:
Cr 2 (SO 4) 3 + K 2 SO 4 + 24H 2 O = 2 + NaCl
4. Redox reactions caused by the properties of the cation or anion. 2KMnO 4 + 16HCl = 2MnCl 2 + 2KCl + 5Cl 2 + 8H 2 O
2. Chemical properties of acid salts:
1. Thermal decomposition with the formation of medium salt
Ca(HCO 3) 2 = CaCO 3 + CO 2 + H 2 O
2. Interaction with alkali. Getting medium salt.
Ba(HCO 3) 2 + Ba(OH) 2 = 2BaCO 3 + 2H 2 O
3. Chemical properties of basic salts:
1. Thermal decomposition. 2 CO 3 = 2CuO + CO 2 + H 2 O
2. Interaction with acid: formation of medium salt.
Sn(OH)Cl + HCl = SnCl 2 + H 2 O
4. Chemical properties of complex salts:
1. Destruction of complexes due to the formation of poorly soluble compounds:
2Cl + K2S = CuS + 2KCl + 4NH3
2. Exchange of ligands between the outer and inner spheres.
K 2 + 6H 2 O = Cl 2 + 2KCl
5.Chemical properties of double salts:
1. Interaction with alkali solutions: KCr(SO 4) 2 + 3KOH = Cr(OH) 3 + 2K 2 SO 4
2. Reduction: KCr(SO 4) 2 + 2H°(Zn, dil. H 2 SO 4) = 2CrSO 4 + H 2 SO 4 + K 2 SO 4
The raw materials for the industrial production of a number of salts - chlorides, sulfates, carbonates, borates Na, K, Ca, Mg are sea and ocean water, natural brines formed during its evaporation, and solid salt deposits. For the group of minerals that form sedimentary salt deposits (sulfates and chlorides of Na, K and Mg), the conventional name “natural salts” is used. The largest deposits of potassium salts are located in Russia (Solikamsk), Canada and Germany, powerful deposits of phosphate ores are in North Africa, Russia and Kazakhstan, NaNO3 is in Chile.
Salts are used in the food, chemical, metallurgical, glass, leather, textile industries, agriculture, medicine, etc.
Main types of salts
1. Borates (oxoborates), salts of boric acids: metaboric HBO 2, orthoboric H3 BO 3 and polyboronic acids not isolated in the free state. Based on the number of boron atoms in the molecule, they are divided into mono-, di, tetra-, hexaborates, etc. Borates are also called by the acids that form them and by the number of moles of B 2 O 3 per 1 mole of the main oxide. Thus, various metaborates can be called monoborates if they contain the B(OH)4 anion or a chain anion (BO2) n n - diborates - if they contain a chain double anion (B 2 O 3 (OH) 2) n 2n- triborates - if they contain a ring anion (B 3 O 6) 3-.
The structures of borates include boron-oxygen groups - “blocks” containing from 1 to 6, and sometimes 9 boron atoms, for example:
The coordination number of boron atoms is 3 (boron-oxygen triangular groups) or 4 (tetrahedral groups). Boron-oxygen groups are the basis of not only island, but also more complex structures - chain, layered and frame polymerized ones. The latter are formed as a result of the elimination of water in hydrated borate molecules and the formation of bridging bonds through oxygen atoms; the process is sometimes accompanied by the cleavage of the B-O bond inside the polyanions. Polyanions can attach side groups - boron-oxygen tetrahedra or triangles, their dimers or extraneous anions.
Ammonium, alkali, as well as other metals in the oxidation state +1 most often form hydrated and anhydrous metaborates such as MBO 2, tetraborates M 2 B 4 O 7, pentaborates MB 5 O 8, as well as decaborates M 4 B 10 O 17 n H 2 O. Alkaline earth and other metals in the oxidation state + 2 usually give hydrated metaborates, triborates M 2 B 6 O 11 and hexaborates MB 6 O 10. as well as anhydrous meta-, ortho- and tetraborates. Metals in the oxidation state + 3 are characterized by hydrated and anhydrous MBO 3 orthoborates.
Borates are colorless amorphous substances or crystals (mainly with a low-symmetric structure - monoclinic or orthorhombic). For anhydrous borates, melting temperatures range from 500 to 2000 °C; The highest melting points are alkali metaborates and ortho- and metaborates of alkaline earth metals. Most borates readily form glasses when their melts are cooled. The hardness of hydrated borates on the Mohs scale is 2-5, anhydrous - up to 9.
Hydrated monoborates lose water of crystallization up to ~180°C, polyborates - at 300-500°C; elimination of water due to OH groups , coordinated around boron atoms occurs up to ~750°C. With complete dehydration, amorphous substances are formed, which at 500-800°C in most cases undergo “borate rearrangement” - crystallization, accompanied (for polyborates) by partial decomposition with the release of B 2 O 3.
Borates of alkali metals, ammonium and T1(I) are soluble in water (especially meta- and pentaborates), and hydrolyze in aqueous solutions (solutions have an alkaline reaction). Most borates are easily decomposed by acids, in some cases by the action of CO 2 ; and SO 2 ;. Borates of alkaline earth and heavy metals interact with solutions of alkalis, carbonates and hydrocarbonates of alkali metals. Anhydrous borates are chemically more stable than hydrated borates. With some alcohols, in particular glycerol, borates form water-soluble complexes. Under the action of strong oxidizing agents, in particular H 2 O 2, or during electrochemical oxidation, borates are converted into peroxoborates .
About 100 natural borates are known, which are mainly salts of Na, Mg, Ca, Fe.
Hydrated borates are obtained: by neutralization of H 3 VO 3 with metal oxides, hydroxides or carbonates; exchange reactions of alkali metal borates, most often Na, with salts of other metals; reaction of mutual transformation of poorly soluble borates with aqueous solutions of alkali metal borates; hydrothermal processes using alkali metal halides as mineralizing additives. Anhydrous borates are obtained by fusion or sintering of B 2 O 3 with metal oxides or carbonates or dehydration of hydrates; Single crystals are grown in solutions of borates in molten oxides, for example Bi 2 O 3.
Borates are used: to obtain other boron compounds; as charge components in the production of glass, glazes, enamels, ceramics; for fire-resistant coatings and impregnations; as components of fluxes for refining, welding and soldering metal”; as pigments and fillers for paints and varnishes; as dyeing mordants, corrosion inhibitors, components of electrolytes, phosphors, etc. Borax and calcium borates are most widely used.
2.Halides, chemical compounds of halogens with other elements. Halides usually include compounds in which the halogen atoms have a greater electronegativity than the other element. Halides are not formed by He, Ne and Ar. To simple or binary EC halides n (n- most often an integer from 1 for monohalides to 7 for IF 7 and ReF 7, but can also be fractional, for example 7/6 for Bi 6 Cl 7) include, in particular, salts of hydrohalic acids and interhalogen compounds (for example, halofluorides). There are also mixed halides, polyhalides, hydrohalides, oxohalides, oxyhalides, hydroxohalides, thiohalides and complex halides. The oxidation number of halogens in halides is usually -1.
Based on the nature of the element-halogen bond, simple halides are divided into ionic and covalent. In reality, the connections are of a mixed nature with a predominance of the contribution of one or another component. Halides of alkali and alkaline earth metals, as well as many mono- and dihalides of other metals, are typical salts in which the ionic nature of the bond predominates. Most of them are relatively refractory, low-volatile, and highly soluble in water; in aqueous solutions almost completely dissociate into ions. Trihalides of rare earth elements also have the properties of salts. The solubility of ionic halides in water generally decreases from iodides to fluorides. Chlorides, bromides and iodides Ag + , Cu + , Hg + and Pb 2+ are poorly soluble in water.
An increase in the number of halogen atoms in metal halides or the ratio of the charge of a metal to the radius of its ion leads to an increase in the covalent component of the bond, a decrease in solubility in water and the thermal stability of halides, an increase in volatility, an increase in oxidation, ability and tendency to hydrolysis. These dependencies are observed for metal halides of the same period and in a series of halides of the same metal. They can be easily observed using the example of thermal properties. For example, for metal halides of the 4th period, the melting and boiling points are respectively 771 and 1430°C for KC1, 772 and 1960°C for CaCl2, 967 and 975°C for ScCl3, -24.1 and 136°C for TiCl4. For UF 3 the melting point is ~ 1500°C, UF 4 1036°C, UF 5 348°C, UF 6 64.0°C. In the rows of connections EH n with constant n The bond covalency usually increases when going from fluorides to chlorides and decreases when going from the latter to bromides and iodides. So, for AlF 3 the sublimation temperature is 1280°C, AlC1 3 180°C, boiling point AlBr 3 254.8°C, AlI 3 407°C. In the series ZrF 4 , ZrCl 4 ZrBr 4 , ZrI 4 the sublimation temperature is 906, 334, 355 and 418°C, respectively. In the ranks of MF n and MC1 n where M is a metal of one subgroup, the covalency of the bond decreases with increasing atomic mass of the metal. There are few metal fluorides and chlorides with approximately equal contributions from the ionic and covalent bond components.
The average element-halogen bond energy decreases when moving from fluorides to iodides and with increasing n(see table).
Many metal halides containing isolated or bridging O atoms (oxo- and oxyhalides, respectively), for example, vanadium oxotrifluoride VOF 3, niobium dioxyfluoride NbO 2 F, tungsten dioxo-iodide WO 2 I 2.
Complex halides (halometallates) contain complex anions in which the halogen atoms are ligands, for example, potassium hexachloroplatinate(IV) K2, sodium heptafluorotantalate(V), Na, lithium hexafluoroarsenate(V). Fluoro-, oxofluoro- and chlorometalates have the greatest thermal stability. By the nature of the bonds, ionic compounds with cations NF 4 +, N 2 F 3 +, C1F 2 +, XeF +, etc. are similar to complex halides.
Many halides are characterized by association and polymerization in the liquid and gas phases with the formation of bridging bonds. The most prone to this are metal halides of groups I and II, AlCl 3, pentafluorides of Sb and transition metals, oxofluorides of the composition MOF 4. Halides with a metal-to-metal bond are known, e.g. Cl-Hg-Hg-Cl.
Fluorides differ significantly in properties from other halides. However, in simple halides these differences are less pronounced than in the halogens themselves, and in complex halides they are less pronounced than in simple halides.
Many covalent halides (especially fluorides) are strong Lewis acids, e.g. AsF 5, SbF 5, BF 3, A1C1 3. Fluorides are part of superacids. Higher halides are reduced by metals and hydrogen, for example:
5WF 6 + W = 6WF 5
TiCl 4 + 2Mg = Ti + 2MgCl 2
UF 6 + H 2 = UF 4 + 2HF
Metal halides of groups V-VIII, except Cr and Mn, are reduced by H 2 to metals, for example:
WF 6 + ZN 2 = W + 6HF
Many covalent and ionic metal halides react with each other to form complex halides, for example:
KS1 + TaCl 5 = K
Lighter halogens can displace heavier halides. Oxygen can oxidize halides, releasing C1 2, Br 2, and I 2. One of the characteristic reactions of covalent halides is interaction with water (hydrolysis) or its vapor when heated (pyrohydrolysis), leading to the formation of oxides, oxy- or oxohalides, hydroxides and hydrogen halides.
Halides are obtained directly from elements, by the reaction of hydrogen halides or hydrohalic acids with elements, oxides, hydroxides or salts, as well as by exchange reactions.
Halides are widely used in technology as starting materials for the production of halogens, alkali and alkaline earth metals, as components of glasses and other inorganic materials; they are intermediate products in the production of rare and some non-ferrous metals, U, Si, Ge, etc.
In nature, halides form separate classes of minerals, which include fluorides (for example, the minerals fluorite, cryolite) and chlorides (sylvite, carnallite). Bromine and iodine are present in some minerals as isomorphic impurities. Significant quantities of halides are contained in sea and ocean water, salt and underground brines. Some halides, for example NaCl, KC1, CaCl 2, are part of living organisms.
3. Carbonates (from Latin carbo, gender carbonis coal), salts of carbonic acid. There are medium carbonates with the CO 3 2- anion and acidic, or hydrocarbonates (old bicarbonates), with the HCO 3 - anion. Carbonates are crystalline substances. Most medium metal salts in the +2 oxidation state crystallize into hexagons. lattice type calcite or rhombic type aragonite.
Of the medium carbonates, only salts of alkali metals, ammonium and Tl(I) are soluble in water. As a result of significant hydrolysis, their solutions have an alkaline reaction. Metal carbonates are most difficult to dissolve in the oxidation state + 2. On the contrary, all bicarbonates are highly soluble in water. During exchange reactions in aqueous solutions between metal salts and Na 2 CO 3, precipitates of medium carbonates are formed in cases where their solubility is significantly less than that of the corresponding hydroxides. This is the case for Ca, Sr and their analogs, the lanthanides, Ag(I), Mn(II), Pb(II) and Cd(II). The remaining cations, when interacting with dissolved carbonates as a result of hydrolysis, can give not intermediate, but basic crabonates or even hydroxides. Medium crabonates containing multiply charged cations can sometimes be precipitated from aqueous solutions in the presence of a large excess of CO 2 .
The chemical properties of carbonates are due to their belonging to the class of inorganic salts of weak acids. The characteristic features of carbonates are associated with their poor solubility, as well as the thermal instability of both the crabonates themselves and H 2 CO 3. These properties are used in the analysis of crabonates, based either on their decomposition with strong acids and the quantitative absorption of the resulting CO 2 by an alkali solution, or on the precipitation of the CO 3 2- ion from solution in the form of BaCO 3. When excess CO 2 acts on a medium carbonate precipitate, hydrogen carbonate is formed in solution, for example: CaCO 3 + H 2 O + CO 2 = Ca(HCO 3) 2. The presence of hydrocarbonates in natural water causes its temporary hardness. Hydrocarbonates, when slightly heated, even at low temperatures, again transform into medium carbonates, which, when heated, decompose to oxide and CO 2. The more active the metal, the higher the decomposition temperature of its carbonate. Thus, Na 2 CO 3 melts without decomposition at 857 °C, and for carbonates Ca, Mg and A1, the equilibrium decomposition pressures reach 0.1 MPa at temperatures of 820, 350 and 100 °C, respectively.
Carbonates are very widespread in nature, which is due to the participation of CO 2 and H 2 O in the processes of mineral formation. carbonates play a large role in global equilibria between gaseous CO 2 in the atmosphere and dissolved CO 2 ;
and HCO 3 - and CO 3 2- ions in the hydrosphere and solid salts in the lithosphere. The most important minerals are calcite CaCO 3, magnesite MgCO 3, siderite FeCO 3, smithsonite ZnCO 3 and some others. Limestone consists mainly of calcite or calcite skeletal remains of organisms, rarely of aragonite. Natural hydrated carbonates of alkali metals and Mg (for example, MgCO 3 ZH 2 O, Na 2 CO 3 10H 2 O), double carbonates [for example, dolomite CaMg(CO 3) 2, trona Na 2 CO 3 NaHCO 3 2H 2 are also known O] and basic [malachite CuCO 3 Cu(OH) 2, hydrocerussite 2PbCO 3 Pb(OH) 2].
The most important are potassium carbonate, calcium carbonate and sodium carbonate. Many natural carbonates are very valuable metal ores (eg carbonates Zn, Fe, Mn, Pb, Cu). Bicarbonates play an important physiological role, being buffer substances that regulate the constancy of blood pH.
4. Nitrates, salts of nitric acid HNO 3. Known for almost all metals; exist both in the form of anhydrous salts M(NO 3) n (n- oxidation state of the metal M), and in the form of crystalline hydrates M(NO 3) n x H 2 O ( X= 1-9). Of aqueous solutions at temperatures close to room temperature, only alkali metal nitrates crystallize as anhydrous, the rest - in the form of crystalline hydrates. The physicochemical properties of anhydrous and hydrated nitrate of the same metal can differ greatly.
Anhydrous crystalline compounds of d-element nitrates are colored. Conventionally, nitrates can be divided into compounds with a predominantly covalent type of bond (salts of Be, Cr, Zn, Fe and other transition metals) and with a predominantly ionic type of bond (salts of alkali and alkaline earth metals). Ionic nitrates are characterized by higher thermal stability, the predominance of crystal structures of higher symmetry (cubic) and the absence of splitting of the nitrate ion bands in the IR spectra. Covalent nitrates have higher solubility in organic solvents, lower thermal stability, and their IR spectra are more complex; Some covalent nitrates are volatile at room temperature, and when dissolved in water, they partially decompose, releasing nitrogen oxides.
All anhydrous nitrates exhibit strong oxidizing properties due to the presence of the NO 3 - ion, while their oxidizing ability increases when moving from ionic to covalent nitrates. The latter decompose in the range of 100-300°C, ionic ones - at 400-600°C (NaNO 3, KNO 3 and some others melt when heated). Decomposition products in solid and liquid phases. are successively nitrites, oxynitrates and oxides, sometimes free metals (when the oxide is unstable, for example Ag 2 O), and in the gas phase - NO, NO 2, O 2 and N 2. The composition of decomposition products depends on the nature of the metal and its degree of oxidation, heating rate, temperature, composition of the gaseous medium, and other conditions. NH 4 NO 3 detonates, and when heated quickly it can decompose with an explosion, in which case N 2, O 2 and H 2 O are formed; when heated slowly, it decomposes into N 2 O and H 2 O.
The free NO 3 - ion in the gas phase has the geometric structure of an equilateral triangle with the N atom in the center, ONO angles ~ 120° and N-O bond lengths of 0.121 nm. In crystalline and gaseous nitrates, the NO 3 - ion mainly retains its shape and size, which determines the space and structure of nitrates. The NO 3 - ion can act as a mono-, bi-, tridentate or bridging ligand, therefore nitrates are characterized by a wide variety of types of crystal structures.
Transition metals in high oxidation states due to steric. Anhydrous nitrates cannot form any difficulties, and they are characterized by oxonitrates, for example UO 2 (NO 3) 2, NbO(NO 3) 3. Nitrates form a large number of double and complex salts with the NO 3 - ion in the internal sphere. In aqueous media, as a result of hydrolysis, transition metal cations form hydroxonitrates (basic nitrates) of variable composition, which can also be isolated in the solid state.
Hydrated nitrates differ from anhydrous nitrates in that in their crystal structures the metal ion is in most cases associated with water molecules rather than with the NO 3 ion. Therefore, they are better soluble in water than anhydrous nitrates, but less soluble in organic solvents; they are weaker oxidizing agents and melt incongruently in water of crystallization in the range of 25-100°C. When hydrated nitrates are heated, anhydrous nitrates, as a rule, are not formed, but thermolysis occurs with the formation of hydroxonitrates and then oxonitrate and metal oxides.
In many of their chemical properties, nitrates are similar to other inorganic salts. The characteristic features of nitrates are due to their very high solubility in water, low thermal stability and the ability to oxidize organic and inorganic compounds. When nitrates are reduced, a mixture of nitrogen-containing products NO 2, NO, N 2 O, N 2 or NH 3 is formed with the predominance of one of them, depending on the type of reducing agent, temperature, reaction of the environment and other factors.
Industrial methods for producing nitrates are based on the absorption of NH 3 by solutions of HNO 3 (for NH 4 NO 3) or on the absorption of nitrous gases (NO + NO 2) by solutions of alkalis or carbonates (for alkali metal nitrates, Ca, Mg, Ba), as well as various exchange reactions of metal salts with HNO 3 or alkali metal nitrates. In the laboratory, to obtain anhydrous nitrates, reactions of transition metals or their compounds with liquid N 2 O 4 and its mixtures with organic solvents or reactions with N 2 O 5 are used.
Nitrates Na, K (sodium and potassium nitrate) are found in the form of natural deposits.
Nitrates are used in many industries. Ammonium nitrite (ammonium nitrate) is the main nitrogen-containing fertilizer; Alkali metal nitrates and Ca are also used as fertilizers. Nitrates are components of rocket fuels, pyrotechnic compositions, etching solutions for dyeing fabrics; They are used for hardening metals, food preservation, as medicines, and for the production of metal oxides.
Nitrates are toxic. They cause pulmonary edema, cough, vomiting, acute cardiovascular failure, etc. The lethal dose of nitrates for humans is 8-15 g, permissible daily intake is 5 mg/kg. For the sum of nitrates Na, K, Ca, NH3 MPC: in water 45 mg/l", in soil 130 mg/kg (hazard class 3); in vegetables and fruits (mg/kg) - potatoes 250, late white cabbage 500, late carrots 250, beets 1400, onions 80, zucchini 400, melons 90, watermelons, grapes, apples, pears 60. Failure to comply with agrotechnical recommendations, excessive application of fertilizers sharply increases the nitrate content in agricultural products, surface runoff from fields ( 40-5500 mg/l), groundwater.
5. Nitrites, salts of nitrous acid HNO 2. Nitrites of alkali metals and ammonium are used primarily, less - alkaline earth and nitrites. d-metals, Pb and Ag. There is only fragmentary information about nitrites of other metals.
Metal nitrites in the +2 oxidation state form crystal hydrates with one, two or four water molecules. Nitrites form double and triple salts, e.g. CsNO 2 AgNO 2 or Ba(NO 2) 2 Ni(NO 2) 2 2KNO 2, as well as complex compounds, for example Na 3.
Crystal structures are known for only a few anhydrous nitrites. The NO 2 anion has a nonlinear configuration; ONO angle 115°, H-O bond length 0.115 nm; the type of M-NO 2 bond is ionic-covalent.
Nitrites K, Na, Ba are well soluble in water, nitrites Ag, Hg, Cu are poorly soluble. With increasing temperature, the solubility of nitrites increases. Almost all nitrites are poorly soluble in alcohols, ethers and low-polar solvents.
Nitrites are thermally unstable; Only nitrites of alkali metals melt without decomposition; nitrites of other metals decompose at 25-300 °C. The mechanism of nitrite decomposition is complex and includes a number of parallel-sequential reactions. The main gaseous decomposition products are NO, NO 2, N 2 and O 2, solid - metal oxide or elemental metal. The release of large amounts of gases causes the explosive decomposition of some nitrites, for example NH 4 NO 2, which decomposes into N 2 and H 2 O.
The characteristic features of nitrites are associated with their thermal instability and the ability of the nitrite ion to be both an oxidizing agent and a reducing agent, depending on the environment and the nature of the reagents. In a neutral environment, nitrites are usually reduced to NO; in an acidic environment, they are oxidized to nitrates. Oxygen and CO 2 do not interact with solid nitrites and their aqueous solutions. Nitrites promote the decomposition of nitrogen-containing organic substances, in particular amines, amides, etc. With organic halides RXH. react to form both nitrites RONO and nitro compounds RNO 2 .
The industrial production of nitrites is based on the absorption of nitrous gas (a mixture of NO + NO 2) with solutions of Na 2 CO 3 or NaOH with sequential crystallization of NaNO 2; Nitrites of other metals are obtained in industry and laboratories by the exchange reaction of metal salts with NaNO 2 or by the reduction of nitrates of these metals.
Nitrites are used for the synthesis of azo dyes, in the production of caprolactam, as oxidizing agents and reducing agents in the rubber, textile and metalworking industries, as food preservatives. Nitrites, such as NaNO 2 and KNO 2, are toxic, causing headaches, vomiting, depressing breathing, etc. When NaNO 2 is poisoned, methemoglobin is formed in the blood and red blood cell membranes are damaged. It is possible to form nitrosamines from NaNO 2 and amines directly in the gastrointestinal tract.
6. Sulfates, salts of sulfuric acid. Medium sulfates with the SO 4 2- anion are known, or hydrosulfates, with the HSO 4 - anion, basic, containing, along with the SO 4 2- anion, OH groups, for example Zn 2 (OH) 2 SO 4. There are also double sulfates containing two different cations. These include two large groups of sulfates - alum , as well as shenites M 2 E (SO 4) 2 6H 2 O , where M is a singly charged cation, E is Mg, Zn and other doubly charged cations. Known triple sulfate K 2 SO 4 MgSO 4 2CaSO 4 2H 2 O (polyhalite mineral), double basic sulfates, for example, minerals of the alunite and jarosite groups M 2 SO 4 Al 2 (SO 4) 3 4Al (OH 3 and M 2 SO 4 Fe 2 (SO 4) 3 4Fe(OH) 3, where M is a singly charged cation. Sulfates can be part of mixed salts, for example 2Na 2 SO 4 Na 2 CO 3 (mineral berkeite), MgSO 4 KCl 3H 2 O (kainite) .
Sulfates are crystalline substances, medium and acidic in most cases, highly soluble in water. Sulfates of calcium, strontium, lead and some others are slightly soluble; BaSO 4 and RaSO 4 are practically insoluble. Basic sulfates are usually poorly soluble or practically insoluble, or are hydrolyzed by water. From aqueous solutions, sulfates can crystallize in the form of crystalline hydrates. Crystal hydrates of some heavy metals are called vitriols; copper sulfate CuSO 4 5H 2 O, iron sulfate FeSO 4 7H 2 O.
Medium alkali metal sulfates are thermally stable, while acid sulfates decompose when heated, turning into pyrosulfates: 2KHSO 4 = H 2 O + K 2 S 2 O 7. Medium sulfates of other metals, as well as basic sulfates, when heated to sufficiently high temperatures, as a rule, decompose with the formation of metal oxides and the release of SO 3.
Sulfates are widely distributed in nature. They are found in the form of minerals, for example, gypsum CaSO 4 H 2 O, mirabilite Na 2 SO 4 10H 2 O, and are also part of sea and river water.
Many sulfates can be obtained by the interaction of H 2 SO 4 with metals, their oxides and hydroxides, as well as the decomposition of volatile acid salts with sulfuric acid.
Inorganic sulfates are widely used. For example, ammonium sulfate is a nitrogen fertilizer, sodium sulfate is used in the glass, paper industries, viscose production, etc. Natural sulfate minerals are raw materials for the industrial production of compounds of various metals, building materials, etc.
7.Sulfites, salts of sulfurous acid H 2 SO 3 . There are medium sulfites with the anion SO 3 2- and acidic (hydrosulfites) with the anion HSO 3 - . Medium sulfites are crystalline substances. Ammonium and alkali metal sulfites are highly soluble in water; solubility (g in 100 g): (NH 4) 2 SO 3 40.0 (13 ° C), K 2 SO 3 106.7 (20 ° C). Hydrosulfites are formed in aqueous solutions. Sulfites of alkaline earth and some other metals are practically insoluble in water; solubility of MgSO 3 1 g in 100 g (40°C). Known crystal hydrates (NH 4) 2 SO 3 H 2 O, Na 2 SO 3 7H 2 O, K 2 SO 3 2H 2 O, MgSO 3 6H 2 O, etc.
Anhydrous sulfites, when heated without access to air in sealed vessels, are disproportionately divided into sulfides and sulfates; when heated in a current of N 2, they lose SO 2, and when heated in air, they are easily oxidized to sulfates. With SO 2 in an aqueous environment, medium sulfites form hydrosulfites. Sulfites are relatively strong reducing agents; they are oxidized in solutions with chlorine, bromine, H 2 O 2, etc. to sulfates. They decompose with strong acids (for example, HC1) with the release of SO 2.
Crystalline hydrosulfites are known for K, Rb, Cs, NH 4 +, they are unstable. The remaining hydrosulfites exist only in aqueous solutions. Density of NH 4 HSO 3 2.03 g/cm 3 ; solubility in water (g in 100 g): NH 4 HSO 3 71.8 (0 ° C), KHSO 3 49 (20 ° C).
When crystalline hydrosulfites Na or K are heated or when the teeming pulp solution is saturated with SO 2 M 2 SO 3, pyrosulfites (obsolete - metabisulfites) M 2 S 2 O 5 are formed - salts of the unknown free pyrosulfuric acid H 2 S 2 O 5; crystals, unstable; density (g/cm3): Na 2 S 2 O 5 1.48, K 2 S 2 O 5 2.34; above ~ 160 °C they decompose with the release of SO 2; dissolve in water (with decomposition to HSO 3 -), solubility (g in 100 g): Na 2 S 2 O 5 64.4, K 2 S 2 O 5 44.7; form hydrates Na 2 S 2 O 5 7H 2 O and ZK 2 S 2 O 5 2H 2 O; reducing agents.
Medium alkali metal sulfites are prepared by reacting an aqueous solution of M 2 CO 3 (or MOH) with SO 2, and MSO 3 by passing SO 2 through an aqueous suspension of MCO 3; They mainly use SO 2 from the exhaust gases of contact sulfuric acid production. Sulfites are used in bleaching, dyeing and printing of fabrics, fibers, leather for grain conservation, green feed, feed industrial waste (NaHSO 3,
Na 2 S 2 O 5). CaSO 3 and Ca(HSO 3) 2 are disinfectants in the winemaking and sugar industries. NaHSO 3, MgSO 3, NH 4 HSO 3 - components of sulfite liquor during pulping; (NH 4) 2 SO 3 - SO 2 absorber; NaHSO 3 is an absorber of H 2 S from industrial waste gases, a reducing agent in the production of sulfur dyes. K 2 S 2 O 5 - a component of acidic fixatives in photography, an antioxidant, an antiseptic.
Methods for separating mixtures
Filtration, separation of heterogeneous systems liquid - solid particles (suspensions) and gas - solid particles using porous filter partitions (FP), which allow liquid or gas to pass through, but retain solid particles. The driving force of the process is the pressure difference on both sides of the phase transition.
When separating suspensions, solid particles usually form a layer of wet sediment on the FP, which, if necessary, is washed with water or other liquid, and also dehydrated by blowing air or other gas through it. Filtration is carried out at a constant pressure difference or at a constant process speed w(the amount of filtrate in m 3 passing through 1 m 2 of the FP surface per unit time). At a constant pressure difference, the suspension is supplied to the filter under vacuum or excess pressure, as well as by a piston pump; When using a centrifugal pump, the pressure difference increases and the process speed decreases.
Depending on the concentration of suspensions, several types of filtration are distinguished. At a concentration of more than 1%, filtration occurs with the formation of a precipitate, and at a concentration of less than 0.1%, with clogging of the pores of the FP (clarification of liquids). If a sufficiently dense layer of sediment does not form on the FP and solid particles enter the filtrate, filter using finely dispersed auxiliary materials (diatomaceous earth, perlite), which are previously applied to the FP or added to the suspension. At an initial concentration of less than 10%, partial separation and thickening of suspensions is possible.
There are continuous and periodic filters. For the latter, the main stages of work are filtering, washing the sediment, its dewatering and unloading. In this case, optimization according to the criteria of greatest productivity and lowest costs is applicable. If washing and dewatering are not carried out, and the hydraulic resistance of the partition can be neglected, then the greatest productivity is achieved when the filtering time is equal to the duration of the auxiliary operations.
Flexible FPs made from cotton, wool, synthetic and glass fabrics are applicable, as well as non-woven FPs made from natural and synthetic fibers and inflexible ones - ceramic, cermet and foam. The directions of movement of the filtrate and the action of gravity can be opposite, coincide or be mutually perpendicular.
Filter designs are varied. One of the most common is a rotating drum vacuum filter (cm. Fig.) of continuous action, in which the directions of movement of the filtrate and the action of gravity are opposite. The distribution device section connects zones I and II with a vacuum source and zones III and IV with a compressed air source. The filtrate and washing liquid from zones I and II enter separate receivers. An automated periodic filter press with horizontal chambers, filter fabric in the form of an endless belt and elastic membranes for dewatering sludge by pressing has also become widespread. It performs alternating operations of filling chambers with suspension, filtering, washing and dewatering sediment, disconnecting adjacent chambers and removing sediment.
