After the start of working days, there was absolutely no time for this matter. So I snatched some time from my family to make some cool stuff.

One of the cool things that I spent half a day with was the Russian motif in copper. He Googles quickly for "Russian ornament vector".

Sources

Free. But in fact, you can just take any Russian painting and vectorize it. I processed it, made the edges not completely black, but checkered, so that the toner would be better transferred. I suffered for a long time and it turned out like this thing

Russian ornament in copper. Etching. Patina.

The main problem with the technology is poor toner transfer. And I sand, and degrease, and warm up, but still there are flaws. For example, below it was normally translated, and on the right, I drew part of the drawing with a marker. There is an idea that it is worth finding a heavy iron.

Attempt to withdraw invoice

Not very happy with the job. She's beautiful, but the untranslated toner ruined everything. And the toner does not translate large polygons, which I do not.
In general, screwing up such a picture is easy. I ruined another engraving photograph, and what I didn’t do I can’t fix. I had to do it again.

ruined work

The work had to be redone. By the way, I found that the concentration of peroxide can be safely reduced. At 2 liters, I have excellent poison with 50 grams. From what I understand, the active ingredient is citric acid.

And the first half of the day yesterday I was busy trying to make a brand to stamp my work, but I didn’t go further than this ...

Stamp blank

The main thing is to transfer the drawing - it is impossible to warm up such an array of steel. To translate from self-adhesives in general is horror. Translating from self-adhesives is just a hell of a gem. Constantly moving out erasing the original drawing. Moved a lot of paper. But the main ambush was with etching. Etched with copper sulphate with salt. The result is so terrible that nafig grind off the result on the skin. In general, it is necessary to poison by the electric method, for this I got an old charger for batteries:

"Velvet"

Or maybe I'll pickle with citric acid.

By the way, I accept orders for drawings in textolite or in copper (a gift for my beloved). Nameplates, etc. We will agree on a price. Not laser engraving, there are flaws, but here it is all the more interesting.

What is especially good about this experience is that everything you need for it is probably at home: a candle, a pharmacy bottle (alcohol solution, iodine tincture) and some worthless iron object - an old door hinge, a key to an unknown lock or a lock whose keys are lost . Grind the metal surface on which the drawing will be with an emery cloth to a shine, light a candle and tilt it so that the paraffin drips onto the shiny surface. Slightly heat the object, then the paraffin will spread in a thin layer. And when it cools and cools down, scratch the grooves with a needle so that they reach the metal. Type pharmacy iodine with a pipette and drop it on scratches. After a few minutes, the iodine solution will turn pale, and then you need to reapply it to the scratches. After about an hour, remove the paraffin layer: you will see clear traces on the metal, they exactly repeat the pattern on the paraffin.

If the experience was successful, you can move on to a more serious task - not just scratching the paraffin, but writing a word or drawing on it, for example, marking your penknife or bicycle wrench.

Let's figure out what happens when iodine comes into contact with metal. Iron reacts with the hearth, resulting in the formation of a salt - iron iodide. And this salt is a powder that is easily removed from the surface. And where there were scratches, depressions formed in the metal. This process is called chemical etching. It is often resorted to, but usually not iodine is used, but other substances that are more active.

By the way, iodine interacts not only with iron, but also with copper. This means that they can pickle various objects made of copper and copper alloys, for example, brass. You can try.

HOME-MADE INDICATORS

In chemical laboratories, indicators are used every now and then - sometimes to determine certain substances, and mostly to find out the acidity of the medium, because both the behavior of substances and the nature of the reaction depend on this property. We will need indicators more than once, and since it is not always possible to buy them, we will try to cook them ourselves. Plants will serve as the initial raw material: many flowers, fruits, berries, leaves and roots contain colored substances that can change their color in response to one or another impact. And, getting into an acidic (or, conversely, alkaline) environment, they clearly signal this to us.

It is not difficult to collect plant "raw materials" in the summer - in the forest, in the field, in the garden or vegetable garden. Take bright flowers - iris, dark tulips and roses, pansies, mallow; dial raspberries, blackberries, blueberries, blueberries; stock up on a few leaves of red cabbage and young beets.

Since indicator solutions are obtained by boiling (a decoction is something like a broth), they, of course, quickly deteriorate - turn sour, moldy. They must be prepared immediately before the experiment. Take some stocked raw materials (the exact amount does not matter), put in a test tube, pour water, put in a water bath and heat until the solution is colored. After cooling, filter each solution and pour into a clean bottle prepared in advance with a label.

To provide yourself with indicators for the whole year, dry the petals and berries in the summer, put them in separate boxes, and then, in the same way as mentioned above, prepare decoctions from them, separately from each plant.

To find out which decoction serves as an indicator for a particular environment and how its color changes, it is necessary to conduct a test. Take a few drops of a homemade indicator with a pipette and add them alternately to an acidic or alkaline solution. Table vinegar can serve as an acidic solution, and a solution of washing soda, sodium carbonate can serve as an alkaline solution. If, for example, you add a bright blue decoction of iris flowers to them, then under the influence of vinegar it will turn red, soda - green-blue.

Record the results of all these experiments carefully, preferably in a table; we present a sample of it here.

Not only leaves and berries can serve you as indicators. Some juices (including those from red cabbage, cherries, black grapes, black currants) and even compotes clearly react to a change in acidity with a change in color. Ordinary borscht can serve as an indicator. Mistresses have noticed this for a long time and use this property of beetroot broth, but not for analysis. To make the borscht bright red, a little food acid - acetic or citric acid - is added to it before the end of cooking; The color changes right before your eyes.

The phenolphthalein indicator is widely used in laboratories. Let's prepare it from pharmaceutical tablets of the same name. Rub one or two tablets and dissolve in about 10 ml of vodka (in extreme cases, just in warm water). In any case, the tablets will not dissolve completely, because in addition to the main substance, phenolphthalein, they also contain a filler - talc or chalk. Filter the resulting solution through blotting paper and pour into a clean bottle labeled "phenolphthalein indicator". This colorless solution does not deteriorate over time. It will come in handy, and more than once, to determine the alkaline environment: in it, it instantly turns red. To test, add a drop or two of phenolphthalein to the washing soda solution.

And here is a sample table that will serve as a guide for you when choosing an indicator:

We suggest you continue the table yourself.

And the last thing about vegetable indicators. It was once fashionable to write invitations on flower petals; but they wrote them, depending on the flower and the desired color of the inscription, with an acid or alkali solution, using a thin pen or a pointed stick. Try, if you like, to write in this way, but choose the petals and writing solutions yourself. Keep in mind that the solution should not be too concentrated, otherwise the delicate petal may be damaged.

EXTRACTION

Now we will get acquainted with a very common process in industry, which is called extraction.

Grind a few nut kernels and a handful of sunflower seeds (of course, without the husk), put in a test tube and fill with gasoline. There should be no fire nearby - gasoline can catch fire! Shake the test tube and let it stand for two hours, remembering to shake it from time to time. Then drain the solution on a saucer and expose to a draft. When the gasoline evaporates, you will see some oil at the bottom. So with the help of gasoline you extracted, extracted, the oil from the seeds. This happened due to the fact that the oil dissolves well in gasoline.

You can try to make oil from other seeds. Just don't hesitate to taste it!

Another experiment is with leaves. For him, we need a water bath and a glass with thin walls (if they are thick, the glass, as you remember, may burst). Place a fresh leaf of some plant in a vessel and fill it with a small amount of diluted alcohol. Heat the water in the bath, remove it from the heat and put a glass with a leaf inside. After some time, remove the leaf with tweezers: it has discolored, and the alcohol has become emerald green. This is how you performed the extraction of chlorophyll, the green pigment in plants.

By the way, if you take a known edible plant - lettuce or spinach, then you can extract food coloring from it in this way - to tint cream or sauce. This is what food factories do: the green edible dye is extracted from the leaves. To speed up this process, we advise you to first grind the leaves and shake the vessel from time to time.

Another experience. In a test tube half filled with water, pour about 1 ml of iodine pharmacy tincture; a brownish solution will be obtained. Add an equal amount of gasoline to it, shake it several times and leave it alone. When the mixture separates, it turns out that the upper, gasoline layer has become dark brown, and the lower, aqueous layer, is almost colorless. Iodine dissolves poorly in water, but well in gasoline. Therefore, he switched from an aqueous solution to a gasoline one.

The difference in solubility is the basis of our latest extraction experience. How to quickly distinguish coffee powder from chicory powder? By smell, this is understandable, but if the smell is weak or you do not remember it exactly? Then throw a pinch of both powders into a transparent vessel with hot water. The colored substances of chicory are difficult to extract with water, so the eye will remain almost colorless. On the contrary, coffee substances dissolve easily in water, and its powder slowly sinks to the bottom, leaving behind a brown trail.

EXPERIMENTS WITH GASES

We have already worked a little with liquids, let's deal with gases. This is somewhat more difficult, and first of all we will need plugs with holes and vent pipes.

The tube can be glass, metal or even plastic. It is better not to take a rubber cork - it is difficult to drill holes in it. Take cork or polyethylene stoppers - holes in them can be burned with a heated awl. Insert a tube into this hole - for example, from an eyedropper; it should enter the cork hole tightly, without gaps, so the hole in the cork must first be made slightly smaller than required, and then gradually expand it, fitting it to the diameter of the tube. Put on a glass tube a rubber or polyethylene flexible tube 30 centimeters long, also insert a short glass tube into its other end.

Now the first experience with gases. Prepare lime water by pouring hot water (1/2 cup) over half a teaspoon of crushed slaked lime, stir the mixture and let stand. A transparent precipitate over the settled solution is lime water. Carefully drain the liquid from the sediment; this laboratory technique, as you remember, is called decantation.

If you do not have slaked lime Ca (OH) 2, then lime water can be prepared from two solutions sold in a pharmacy: calcium chloride CaCl 2 and ammonia NH 4 OH (ammonia aqueous solution). When mixed together, clear lime water is also obtained.

Grab a chilled bottle of mineral water or lemonade. Open the cork, quickly insert the cork with the gas outlet tube into the neck, and lower its other end into a glass of lime water. Place the bottle in warm water. Gas bubbles will come out of it. This is carbon dioxide CO 2 (aka carbon dioxide, carbon dioxide). It is added to water to make it tastier.

The gas enters the glass through the tube, it passes through lime water and it becomes cloudy before our eyes, because the calcium hydroxide contained in it turns into calcium carbonate CaCO 3, and it is poorly soluble in water and forms a white turbidity.

To experiment with lime water, it is not necessary to buy lemonade or mineral water. After all, when we breathe, we consume oxygen and release carbon dioxide, the same gas that makes lime water cloudy. Dip the end of any clean tube into a fresh portion of lime water and exhale through the tube several times - the result will not be long in coming.

Open another bottle, insert a stopper with a tube and continue to pass carbon dioxide through the lime water. Some time later, the solution will again become transparent, because carbon dioxide reacts with calcium carbonate, turning it into another salt - bicarbonate Ca (HCO 3) 2, and this salt just dissolves very well in water.

The next gas we'll be looking at has been mentioned quite recently: ammonia. It is easy to recognize by its sharp characteristic smell - the smell of pharmacy ammonia.

Pour some boiled saturated washing soda solution into the bottle. Then add ammonia, insert a stopper with a flexible outlet tube into the neck and put the test tube upside down on its other end. Warm the bottle in warm water. Ammonia vapor is lighter than air and will soon fill an inverted test tube. Still holding the tube upside down, carefully lower it into the beaker of water. Almost immediately, the water will begin to rise up into the test tube, because ammonia is highly soluble in water, making room for it in the test tube.

At the same time, you can learn to recognize ammonia - and not just by smell. First, make sure the ammonia solution is alkaline (use phenolphthalein or homemade indicators). And secondly, conduct a qualitative reaction for ammonia. A qualitative reaction is one that allows you to accurately identify a particular substance or group of substances.

Prepare a weak solution of copper sulphate (it should be pale blue) and lower the gas outlet tube into it. When ammonia NH 3 begins to be released, the solution will turn bright blue at the end of the tube. Ammonia with a copper salt gives a brightly colored complex compound of a rather complex SO 4 composition.

Now try to get a very small piece of calcium carbide - we will get acetylene. Assemble the device, as in the previous experiment, only pour not ammonia into the bottle, but soda. Dip a small, pea-sized piece of calcium carbide, carefully wrapped in blotting paper, into it and insert a cork with a tube. When the blotting paper gets wet, gas will begin to be released, which you will collect in an inverted test tube as before. After a minute, turn the test tube upside down and bring a lit match. The gas will ignite and burn with a smoky flame. This is the same acetylene that gas welders use.

By the way, not only acetylene is obtained in this experiment. An aqueous solution of calcium hydroxide, i.e. lime water, remains in the bottle. It can be used for experiments with carbon dioxide.

The next experiment with gases can be done only with good ventilation, and if it is not, then in the fresh air. We will get a sharp smelling sulfur dioxide (sulphurous gas) SO 2 .

Pour dilute acetic acid into a bottle and add some sodium sulfite Na 2 SO 3 wrapped in blotting paper (this substance is sold in photo stores). Close the bottle with a cork, lower the free end of the gas outlet tube into a glass with a dilute solution of potassium permanganate KMnO 4 prepared in advance (this substance is known in everyday life as potassium permanganate). The solution should be pale pink. When the paper gets wet, sulfur dioxide will start to come out of the bottle. It reacts with potassium permanganate solution and discolors it.

If you can't buy sodium sulfite, then replace it with the contents of a large cartridge of a conventional photodeveloper. True, in this case there will be an admixture of carbon dioxide in sulfur dioxide, but this will not interfere with the experiment.

OXIDATION-REDUCTION

The experience with sulfur dioxide showed us one of the many redox reactions. In such reactions, atoms of some substances gain electrons, while others donate electrons. The first are called oxidizing agents (potassium permanganate), the second - reducing agents (sulfur dioxide).

Let's put some more experiments with oxidation - reduction.

On a fresh slice of potato, drip dilute iodine tincture: a blue color will appear. It is the starch found in potatoes that turns blue in the presence of free iodine. Such a reaction is often used to detect starch, which means that this is also a qualitative reaction.

In the same place where you dropped the iodine tincture, pour a little sodium sulfite solution. The color will fade quickly. This is what happened: sulfite gave an electron to free iodine, it became electrically charged, turned into an ion, and in this state, iodine no longer reacts with starch.

This property of sodium sulfite, like sulfur dioxide, means that these substances are good reducing agents. Here is another interesting experience with sulfite. Its oxidizing companion will again be potassium permanganate.

Pour pale pink, pink, light purple and dark purple solutions of potassium permanganate into four test tubes. Add sodium sulfite solution to each tube. The contents of the first tube will become almost colorless, the second - brownish. In the third tube, brown flakes will fall out, in the fourth, too, but the sediment will be much larger. In all test tubes, solid manganese oxide MnO 2 is formed. But in the first two test tubes, it exists as a colloidal solution (the solid particles are so small that the solution appears clear). And in the remaining two test tubes, the concentration of MnO 2 is so high that the particles stick together and precipitate.

In general, potassium permanganate resembles a chemical chameleon - this is how it can change its color. For example, in an alkaline environment, a solution of potassium permanganate turns from red-violet to green because the permanganate is reduced to green manganate. To check this, drop a potassium permanganate crystal into an alkali solution - into a concentrated boiled solution of washing soda, and green will appear instead of the usual pink color.

This experience is even more beautiful when working with caustic soda, but for home experimentation, until you have the skill and ability, such alkalis cannot be recommended. If you are studying in a circle, then set up the experiment like this: pour a little red solution of potassium permanganate into a thin-walled glass (it should be transparent) and in very small portions so that the reaction mixture does not warm up, add a sufficiently concentrated solution of caustic soda. Watch the color of the liquid - first it will become more and more purple, then, as the alkalinity increases, blue, and finally green.

The color change is especially clearly visible in transmitted light. In any case, the lighting should be good, without this, the transitions of shades may not be noticed.

The following experiment will help you distinguish dirty water from clean water. Fill one test tube with clean water, the other with water from a stagnant puddle or from a swamp. Add a little solution of the oxidizing agent - potassium permanganate to the test tubes. In tap water, it will remain pink; in puddle water, it will become discolored. In warm weather, organic matter accumulates in stagnant water. They, like sodium sulfite, restore potassium permanganate, change its color.

In the first experiment with sodium sulfite, it was proposed to take it from a large developer cartridge. If you followed this advice, then you are left with a small cartridge that contains a mixture of metol and hydroquinone. Dissolve this mixture in water; the solution will be very faintly colored. Add some bleach (it's a common disinfectant and should be handled with care). The contents of the tube will turn yellow. Chlorine is a good oxidizing agent, it oxidizes hydroquinone to quinone, which is colored yellow. If, however, a mixture of sodium sulfite and soda is now added to the test tube from a large cartridge, the yellow color will disappear: sodium sulfite will again reduce quinone to hydroquinone.

We will deliver the last experiment on the topic "oxidation - reduction" with chromium compounds. Such experiences are often colorful, which is not surprising, since “lame” in Greek means “color”.

So, take some yellow solution of potassium dichromate K 2 Cr 2 O 7; this substance is widely used in engineering as an oxidizing agent, for example, for cleaning heavily contaminated parts; it must be handled with care. If you add a little sulfuric acid to the yellow solution (carefully! pour the acid slowly!), then it will turn red. Throw a few pieces of zinc into such an acidified solution. If you do not have granular zinc, which is usually used for experiments, then extract zinc yourself from an unusable battery: metal cups in batteries are zinc.

So, you threw a little zinc into a glass with a solution, and the bichromate, recovering, changes color to dark green. It formed Cr 3+ ions. At the same time, due to the reaction of zinc with acid, gas is released - hydrogen. If the reaction products are not oxidized by atmospheric oxygen, then the reaction will continue, and a blue color will appear - this is the color of a solution of chromium sulfate CrSO 4. Pour it into another glass; while you do this, oxidation will occur and the solution will turn green again.

ADSORPTION

Everyone is probably familiar with the physicochemical phenomenon that will be discussed now, although, perhaps, not everyone knows that it is called adsorption. Even if you did not go through adsorption in the classroom, you observed it more than once. As soon as you plant an ink blot on paper or, much worse, on clothes, you immediately get acquainted with this phenomenon. When the surface of one substance (paper, cloth, etc.) absorbs particles of another substance (ink, etc.), this is adsorption.

Very good adsorbent - coal. And not stone, but woody, and not just woody, but active (activated). Such coal is sold in pharmacies, usually in the form of tablets. We will begin experiments on adsorption with it.

Prepare a pale ink solution of any color and pour into a test tube, but not to the top. Put a tablet of activated charcoal, preferably crushed, into a test tube, close with your finger and shake well. The solution will brighten before your eyes. Change the solution to some other, but also colored one - let it be diluted gouache or watercolor. The effect will be the same. And if you just take pieces of charcoal, they will absorb the dye much weaker.

There is nothing strange in this: activated carbon differs from ordinary carbon in that it has a much larger surface. Its particles are literally permeated with pores (for this, coal is processed in a special way and removed from its impurities). And since adsorption is absorption by the surface, it is clear: the larger the surface, the better the absorption. Adsorbents are able to absorb substances not only from solutions. Take a half-liter glass jar and put one drop of cologne or any other odorous substance on the bottom. Grab the jar with your palms and hold it like that for half a minute to heat the odorous liquid a little - then it will evaporate faster and smell stronger. As is customary in chemistry, do not sniff the substance directly from the bottle, but with a slight wave of the hand direct the air along with the vapors of the substance to the nose; it is not always known whether the substance in the bottle smells good.

Whatever the smell, you will certainly feel it clearly. Now put some activated charcoal in the bottle, close it tightly with a lid and leave for a few minutes. Remove the lid and again direct the air towards you with a wave of your palm. The smell is gone. It was absorbed by the adsorbent, or, more precisely, the molecules of the volatile substance that you placed in the jar were absorbed.

It is not necessary to take active carbon for these experiments. There are many other substances that can serve as adsorbents: tuff, dry ground clay, chalk, blotting paper. In a word, a variety of substances, but always with a developed surface. Including some food products - you probably know how easily bread absorbs odors. It is not for nothing that wheat bread is not advised to be kept in the same package with rye bread - their smells are mixed, and each one loses its special, unique aroma.

A very good adsorbent is puffed corn, or corn sticks, so loved by many of us. Of course, it makes no sense to spend a package or even a quarter of a package on experience, but a few pieces ... Let's try. Repeat the previous experiment with odorous substances in the presence of corn sticks - and the smell will completely disappear. Of course, after the experience, it is no longer possible to eat sticks.

Let's return to the experiment with the production of carbon dioxide (carbon dioxide). Fill two test tubes with this gas, put corn sticks in one and shake several times. Then, as before, do the experiment with lime water (you can simply “pour” gas from test tubes into it - it is heavier than air). Will there be a difference in the behavior of lime water? Yes, it will. The liquid will become cloudy only in the glass into which the gas, which has not been treated with an adsorbent, has been “poured”. And from another test tube, the one where there were corn sticks, carbon dioxide cannot be extracted: it was absorbed by the adsorbent.

If you work in a chemistry circle and have already learned how to obtain and collect colored gases such as chlorine and nitric oxide (you don’t need to deal with them at home, good traction is required here), then you can test the effect of coal and corn sticks on them. Place an adsorbent into a vessel with a colored gas, shake it several times - and the color, if it does not disappear completely, will noticeably weaken.

Now in many kitchens, a variety of devices are placed above gas stoves to clean the air from fumes and smoke. In such devices, among other things, there is a cartridge with some kind of adsorbent through which polluted air is driven. What happens, you now know. And when the entire surface is occupied by foreign particles “absorbed” from the air, the cartridge is replaced with a fresh one.

DRY CLEANING

The experiments in this chapter can be called a repetition of the past, because dry cleaning and stain removal most often use exactly the same processes that you recently met in the experiments. Namely: extraction, oxidation - reduction and adsorption.

Of course, it is not worth getting your clothes dirty for the sake of experiments. Let's do this: prepare a few pieces of light fabric, put different spots on it and try to remove them. And if the experiments are successful, you can risk cleaning your own suit (or someone else's - if allowed ...).

The most common stains are fat. They are removed, as a rule, by means of extraction, selecting a suitable solvent for this. To remove fresh grease stains, gasoline, turpentine, and medical ether are suitable. With a cotton swab moistened with a solvent, wipe the stain several times, and the fat will go into solution. So that there is no halo on the fabric, it must be wiped with soapy water or a solution of washing powder.

Old grease stains are more difficult to remove, here one cannot do with a solvent, mixtures are needed. For example, gasoline, medical ether and turpentine (7:1:2) or wine spirit, turpentine and medical ether (10:2:1).

If the fabric is colored, then care must be taken that the solvent does not damage the color. Before you begin, check to see if the thinner you choose does not change the color of the fabric.

A stain from oil varnish is well removed by a paste of gasoline and white clay. The paste-like mixture is applied to the stain and left until the gasoline has completely evaporated. In this case, adsorption is added to the extraction: white clay absorbs, absorbs substances extracted by gasoline.

Moisten a fresh stain from oil paint first with turpentine (to soften), and then remove with gasoline. If such treatment may damage the color, then wipe the stain with a hot solution of glycerin or its mixture with an equal amount of ethyl alcohol.

Extraction can also remove grass stains. Remember the experience in which we extracted chlorophyll with alcohol? So, if you wipe the stained place with alcohol (or medical ether), you can gradually extract chlorophyll from the stain, and it will discolor.

Ink stains planted on clothes can also sometimes be discolored. To do this, sprinkle a little crushed chalk or tooth powder on the stain and drip 2-3 drops of alcohol. The alcohol will dissolve the ink dye, and the chalk will absorb the colored solution. Remove the soiled chalk with the blunt end of a knife, apply fresh chalk and alcohol, and repeat this operation until the chalk remains white. Let it dry and remove the residue with a brush.

And in this case, we combined extraction with adsorption. In general, when removing stains, this double technique is often the most effective: white clay, chalk, and similar powders do not allow the tinted solution to spread over the fabric, forming a halo around the former stain.

Now about redox reactions, which also help remove stains.

Fresh stains from berries and juices can often be removed simply with hot water. If this does not work, then these spots on white fabrics can be discolored with a solution of hydrogen peroxide (you can dissolve a tablet of hydroperite in half a glass of water). Saturate the stain with this solution, adding a few drops of ammonia to it, wipe it with a clean cotton swab and rinse with water. Hydrogen peroxide (peroxide) is a strong oxidizing agent, it oxidizes many dyes, and they discolor.

Hot iron stains on white cotton and linen fabrics can also be removed using an oxidation-reduction reaction. An aqueous solution of bleach should be used as an oxidizing agent (carefully!) in a ratio of 1:50 by weight. When the tissue is overheated, brown products of thermal oxidation are formed, and bleach destroys them, making them colorless. But keep in mind that the reaction produces hydrochloric (hydrochloric) acid, which itself can destroy tissue. Therefore, immediately after cleaning, rinse the fabric with a weak solution of soda to neutralize the acid, and then rinse with clean water.

Finally, if iodine has got on the fabric, then by wiping the stain with a solution of sodium thiosulfate (hyposulfite), you will remove the stain without a trace. You already know what is an oxidizing agent and what is a reducing agent in this reaction.

From dry cleaning it will be quite natural to move on to laundry, which we will do.

Washing is a physical and chemical process, its main actors are surfactants. The molecules of such substances consist of two parts - hydrophilic, that is, having an affinity for water, and hydrophobic, which does not interact with water, but willingly comes into contact with pollutants, for example, with difficult-to-wash fats and oils. These groups - hydrophilic and hydrophobic - are located at different ends of a long molecule. Such molecules attach their hydrophobic ends to the oily surface, while the hydrophilic ones stick out like needles on a hedgehog. Water wets these "needles" well, it surrounds such a "hedgehog", tears it off the surface and carries it away. This is how soap and laundry detergent work. And in order to quickly remove dirt from the fabric or from our hands, we rub them with a sponge, a brush, against each other ...

Since soap is the oldest surfactant, let's start with it.

Dissolve a little soap in a small amount of water, add a solution of phenolphthalein to the test tube. The color will turn crimson red. So the environment is alkaline. Indeed, ordinary soap is the sodium salt of fatty acids - oleic, stearic, for example, C 17 H 35 COONa (and liquid soap is the potassium salt of the same acids). When dissolved in water, such salts hydrolyze, decompose into an acid and an alkali. But fatty acids are weak, and alkalis in this case are strong, so the solution has an alkaline reaction.

It used to be thought that soap washes and washes well because it forms alkali. It turned out that this was not the case at all. In contrast, an alkali (like washing soda) cleans because it combines with fats and forms soap-like surfactants in solution.

By the way, soap is not that hard to get yourself. There are several ways; here is one of them. Prepare a hot concentrated solution of washing soda, pour it into a test tube and gradually, drop by drop, add vegetable oil until it no longer dissolves. Instead of oil, you can take beeswax. Pour a pinch of table salt into the resulting solution. So do soap factories - this process is called salting out. After adding salt, solid soap floats to the surface and is easily separated from the solution.

Now soap for washing is used less and less, and washing powders are increasingly used. The composition of these powders includes surfactants obtained synthetically. That is why they are called synthetic detergents.

Let's put this experience. Cut a piece of dirty cloth into three pieces and dip each piece into glasses. Pour just warmed water into the first glass, a soap solution into the second, and a solution of any washing powder that you can find at home into the third. Lightly rub the patches, rinse them in clean water, dry and carefully examine. That piece of cloth that had been in the water was not much cleaner. A piece of soapy solution brightened noticeably. But the cleanest will be the piece of fabric that you removed from the glass with a solution of washing powder. This means that synthetic detergents act more vigorously than ordinary soap.

Many washing powders have another valuable property: they wash in any water - soft, hard, even sea water. What about soap?

Take ordinary water and dissolve some calcium or magnesium salt in it. You can buy bitter salt at the pharmacy, you can take dry sea salt (it is also sold in pharmacies) or a solution of calcium chloride. In this way, you will make the water hard, because hard water differs from soft water in that it contains a lot of calcium and magnesium salts - the so-called hardness salts.

Again, take a piece of dirty cloth and try washing it with soap in such hard water. Nothing will come of it - not even foam is formed. Hardness salts react with soap, calcium and magnesium soaps are formed, and they are insoluble in water. And our soap loses all its useful properties.

But if you dissolve washing powder in hard water, for example, Lotus, it will wash dirt almost the same as before - hard water does not harm it. The surfactants that make up the powder do not interact with hardness salts, and therefore do not lose their properties.

Solutions of washing powders, as well as solutions of laundry soap, can be alkaline; in this case, they recommend washing cotton and linen, but not wool or silk. However, there are also neutral agents, they are often produced not in the form of powders, but in the form of liquids; they are good for wool, silk and synthetic fabrics. If in doubt whether it is worth washing a wool sweater with one or another powder, then test with phenolphthalein. The solution turned red - it means that there is free alkali in it, and it is contraindicated for wool, because it can destroy the fibers. But if the solution remains colorless or stained quite a bit, feel free to immerse both woolen and silk items in it.

In the old days, when soap was a luxury item, other, more affordable substances were often used for washing, which, although to a lesser extent, still washed off the dirt. Try and you, how these substances work. You can take mustard powder or a decoction of beans for the experiment, but even better - the roots of some plants, for example, primrose, crow's eye, cyclamen, cockle. These roots contain saponins - substances that have a washing effect (perhaps in old books you have come across such an expression - soap root). All these natural substances wash, of course, worse than soap, but you can easily make sure that they still wash.

Let's finish the chapter on detergents with an experiment in which, by adding surfactants and thereby changing the surface tension of water, we will make an object move through the water.

From a thin copper wire, make a flat spiral in several turns, lightly grease it with oil or petroleum jelly and lower it very carefully to the surface of the water. The surface tension of water prevents the spiral from sinking, and water does not wet it. Now, with a pipette, gently drop one drop of soapy water into the very middle of the spiral. The spiral will now begin to spin. Spreading over the surface, the soap solution reaches the end of the spiral, leaves it and develops a small jet thrust. When the spiral stops, drip the soap solution again - the rotation will resume.

Such a spiral can serve as a device for determining the surface activity of various liquids. Replace the soap solution with another substance - the spiral will move at a different speed. If you drop a solution of table salt, then there will be no circular motion at all. And in a solution of washing powder, the spiral will quickly sink. It washes away the layer of oil that holds the wire in the water.

SOAP CANDLE

When we talked about why soap washes, we mentioned the special structure of its molecule: a “head” and a long “tail”, moreover, the “head” tends to water, and the “tail”, on the contrary, is repelled from water ...

Let's take a closer look at the hydrophobic "tail" - a long hydrocarbon chain. Such compounds are very common and extremely important for industry. They are an indispensable component of many fats, oils, lubricants and other useful substances. One of them - the so-called stearin - we will now get, taking laundry soap as a basis.

With a knife, cut from half a piece of laundry soap and put in a clean tin can (or in your used saucepan). Pour in enough water to cover the soap chips, and place the mixture in a water bath. Stir the contents of the saucepan with a wooden stick from time to time so that the soap quickly dissolves in the water. When this finally happens, remove the vessel from the fire (of course, not with your bare hand) and pour vinegar into it. Under the action of acid, a thick white mass will stand out from the solution and float to the surface. This is stearin - a translucent mixture of several substances, mainly stearic C 17 H 35 COOH and palmitic C 15 H 31 COOH acids. It is impossible to say the exact composition, it depends on the substances that went into the preparation of soap.

From stearin, as is known from fiction, candles are made. Or rather, they did it before, because now candles are mostly not stearin, but paraffin - paraffin obtained from oil is cheaper and more accessible. But, as soon as we have stearin at our disposal, we will prepare a candle from it. That, by the way, is a fun activity in itself!

When the jar is completely cool, scoop the stearin off the surface with a spoon and transfer it to a clean bowl. Rinse the stearin two or three times with water and wrap in a clean white cloth or filter paper to absorb excess moisture. When the stearin is completely dry, let's start with a candle.

Here is perhaps the simplest trick: dip a thick twisted thread, for example, from a wick for a kerosene stove, repeatedly into slightly heated melted stearin, each time allowing the stearin to harden on the wick. Proceed in this way until a candle of sufficient thickness grows on the wick. This is a good way, although somewhat tedious; in any case, in ancient times, candles were often prepared in this way.

There is a simpler way: immediately coat the wick with stearin heated to soften (it can even be freshly cooked, not yet cooled down). But in this case, the wick will be worse saturated with fusible mass and the candle will turn out not very good, although it will burn.

For beautiful, figured candles, manufacturing methods are not easy. And first of all, you need to make a form - wooden, plaster, metal. In this case, too, it is desirable to impregnate the wick first with one or two layers of stearin; then it is fixed in the form so that it passes exactly in the middle. It is desirable that the wick be slightly taut. And after that, hot stearin is poured into the mold.

By the way, in this way you can make candles from paraffin, i.e., in fact, from purchased candles, by melting them and giving them the shape you like. However, we warn you - you will have to tinker ...

Having received a candle from soap, we will carry out the experiment in the opposite direction: we will prepare soap from a candle. Only not from paraffin, soap cannot be made from it at all, because paraffin molecules do not have “heads”. But if you are sure that the candle is stearic, then you can safely prepare laundry soap from it. Natural beeswax is also suitable.

Heat a few fragments of a stearin candle in a water bath, hot enough, but not brought to a boil. When the stearin is completely melted, add a concentrated solution of washing (soda ash) to it. The resulting white viscous mass is soap. Hold it for a few more minutes in a water bath, and then, putting on a mitten or wrapping your hand in a towel so as not to burn yourself, pour the still hot mass into any form - at least into a matchbox. When the soap has hardened, take it out of the box.

To make sure that this is soap and that it washes will not be difficult. Just please don't use it to wash your hands, because we don't know how pure the ingredients in the candle were.

CHALK, MARBLE, SHELL…

Moisten a piece of natural chalk CaCO 3 with a drop of hydrochloric acid Hcl (you can take pharmacy acid). Where the drop fell, a vigorous effervescence is noticeable. Put a piece of chalk with a “boiling” drop into the flame of a candle or dry alcohol. The flame will turn into a beautiful red color.

This phenomenon is well-known: calcium, which is part of the chalk, makes the flame red. But why acid? It, reacting with chalk, forms soluble calcium chloride CaCl 2, its splashes are carried away by gases and fall directly into the flame - this makes the experience more effective.

Unfortunately, such an experiment with pressed school chalk fails - it contains an admixture of soda (sodium salts), and the flame turns orange. The best experience is obtained with a piece of white marble moistened with the same acid. And you can make sure that sodium salts color the flame in an intense yellow color by adding a grain of NaCl salt to the flame (or just lightly “salting” the fire).

For the next experiment with chalk, you will need a candle. Mount it on a non-combustible stand and add a piece of chalk (marble, shells, eggshell) to the flame. The chalk is covered with soot, which means that the temperature of the flame is low. We are going to burn the chalk, and for this we need a temperature of 700-800 ° C. How to be? It is necessary to increase the temperature by blowing air through the flame.

Remove the rubber cap from the pharmacy pipette and put on a rubber or plastic tube instead. Blow into the tube in such a way that air enters the flame just above the wick through the retracted end of the pipette. The tongue of flame will deviate to the side, its temperature will rise. Point the tongue at the sharpest part of the crayon. This area will become white hot, the chalk will turn into burnt (quicklime) lime CaO, and at the same time carbon dioxide will be released.

Do this operation several times with pieces of chalk, marble, eggshells. Put the burnt pieces into a clean tin. While they are cooling, place the largest piece in a saucer and drop water on the place that was heated. There will be a hiss, all the water will be absorbed, and the calcined area will crumble into powder. This powder is Ca (OH) 2 slaked lime.

Add more water and drip a solution of phenolphthalein. The water in the saucer will turn red; this means that slaked lime forms an alkaline solution.

When the burned pieces have cooled, place them in a glass jar or bottle, fill with water, close the lid and shake - the water will become cloudy. You already know that we will now receive lime water. Let the liquid settle and pour the clear solution into a clean bottle. Pour some lime water into a test tube - and you can use it to perform the previously described experiments with gases. And you can do tricks, like turning “water” into “milk” or “water” into “blood”. You will find a description of such tricks in the "Sleight of Hand" section.

ELECTROLYSIS IN A GLASS

Experiments with electricity will meet you more than once in this book. Now - the most simple. To conduct them, three or four batteries for a flashlight are enough.

In fact, experiments in electrochemistry are often tried at home, but they do not always come out: some little thing - and nothing happens. If you follow all our instructions, you can be sure that the experience will succeed.

Let's start with a very simple yet instructive experience. For him, you need one single reagent: ink of any color. True, you have to work a little on the device.

Take two metal strips 8-10 cm long and 1-2 cm wide. They can be made of iron, copper, aluminum - it doesn't matter, as long as they freely pass into a transparent vessel - a tall beaker or a large test tube. Before experiment, drill holes in the plates on one side for attaching conductors. Prepare two identical, literally a few millimeters thick, plastic or wooden spacers and glue them with metal strips so that they, arranged in parallel, do not touch each other. Almost any glue is suitable - BF, Moment, etc.

Pour water into a beaker or test tube and drop enough ink into it so that the solution is not very saturated in color (however, it should not be transparent either). Lower the construction of two strips into it, connect them with wires to two batteries connected in series, “plus” to “minus”. A few minutes later, the ink solution between the plates will begin to lighten, and dark particles will collect at the bottom and top.

The composition of the ink consists of very small colored particles suspended in water. Under the action of current, they stick together and can no longer swim in the water, but sink to the bottom under the action of gravity. It is clear that the solution thus becomes more and more pale.

But how did the particles get to the top? When a current is applied to solutions, gases are often formed. In our case, gas bubbles pick up solid particles and carry them up.

In the next experiment, a thick-walled tea glass, expanding upwards, will serve as an electrolytic bath. Prepare a plywood circle of such a diameter that it is pressed against the wall of the glass three to four centimeters above the bottom. Drill two holes in the circle in advance (or cut a slot in it in diameter), pierce two holes nearby with an awl: the wires will pass through them. Insert two 5-6 cm long pencils, sharpened at one end, into large holes or into a slot. Pencils, more precisely, their leads, will serve as electrodes. On the unfinished ends of the pencils, make notches so that the leads are exposed, and attach the bare ends of the wires to them. Twist the wires and carefully wrap them with insulating tape, and in order for the insulation to be completely reliable, it is best to hide the wires in rubber tubes. All parts of the device are ready, it remains only to assemble it, that is, insert a circle with electrodes inside the glass.

Place the glass on a plate and fill it to the brim with a solution of washing soda ash Na 2 CO 3 at the rate of 2-3 teaspoons per glass of water. Fill two test tubes with the same solution. Close one of them with your thumb, turn it upside down and immerse it in a glass so that not a single air bubble gets into it. Under water, put the test tube on the pencil electrode. Do the same with the second tube.

Batteries - at least three in number - must be connected in series, "plus" one to the "minus" of the other, and connect wires from pencils to the extreme batteries. The electrolysis of the solution will start immediately. Positively charged hydrogen ions H + will go to the negatively charged electrode - the cathode, attach an electron there and turn into hydrogen gas. When a full tube of hydrogen is gathered at the pencil connected to the “minus”, it can be removed and, without turning over, set fire to the gas. It will light up with a characteristic sound. At the other electrode, positive (anode), oxygen is released. Close the test tube filled with it with your finger under water, remove it from the glass, turn it over and bring in a smoldering splinter - it will light up.

So, from water H 2 O, both hydrogen H 2 and oxygen O 2 were obtained; what is the soda for? To speed up the experience. Pure water conducts electricity very poorly, the electrochemical reaction in it is too slow.

With the same device, you can put another experiment - the electrolysis of a saturated sodium chloride solution NaCl. In this case, one tube will be filled with colorless hydrogen, and the other with yellow-green gas. This is chlorine, which is formed from table salt. Chlorine easily gives up its charge and is the first to be released at the anode.

Close the test tube with chlorine, which also contains a little salt solution, with your finger under water, turn over and shake without removing your finger. In a test tube, a solution of chlorine is formed - chlorine water. It has strong bleaching properties. For example, if you add chlorine water to a pale blue ink solution, it will discolor.

During the electrolysis of table salt, another substance is formed - caustic soda. This alkali remains in solution, which can be verified by dropping a little phenolphthalein solution or a homemade indicator into a glass near the negative electrode.

So, in the experiment we got three valuable substances at once - hydrogen, chlorine and caustic soda. That is why salt electrolysis is so widely used in industry.

With the help of current and a saturated solution of sodium chloride, another interesting experiment can be done. Let's now deal with the fact that we will drill metal with an ordinary pencil.

Prepare a saturated salt solution in a tea saucer. Wire a safety razor blade to the positive pole of a flashlight battery (the blade will be the anode). At the sharpened end of the pencil, break off the stylus and pick it out about half a millimeter with a needle. 2–3 cm higher, make a notch with a knife to the stylus and wind the end of the bare wire around it; wrap this place with insulating tape, and attach the other end of the wire to the negative pole of the battery (the pencil will be the cathode).

Place the blade in a saucer of solution and touch the cathode pencil to the blade. Immediately, hydrogen bubbles will begin to bubble up around the pencil. And the anode blade will dissolve: the iron atoms will acquire a charge, turn into ions and go into solution. So in ten to fifteen minutes a through hole will turn out in the blade. It forms especially quickly if the battery is new and the blade is thin (0.08 mm). In aluminum foil, a hole is drilled in just seconds.

If you want to drill a hole with a pencil in a certain place on a thin metal plate, then it is better to varnish the workpiece in advance, and remove the varnish where you will drill.

The recess in the stylus was then needed so that the stylus did not touch the metal. Otherwise, the circuit will immediately close, the current will not go through the solution and there will be no electrolysis.

You can drill with a pencil without an electrolytic bath (in our case, without a tea saucer). Put the anode plate on a board or on a plate, drip water, dip the pencil attached to the battery in salt and immerse its sharpened end in a drop. From time to time, remove the electrolysis products with a cloth and apply a new drop. By repeating this operation, you can, without effort, drill through metal foil or tin from a tin can. Also, by the way, you can make a hole in a broken steel knife to attach a new handle to it.

Of course, for drilling metal with a thickness of more than a millimeter, one battery is not enough - you need to connect several batteries in parallel or use a step-down transformer with a rectifier - for example, from a children's railway or from a wood burning device. And regardless of the current source and method of electrolysis, you will have to change the electrolyte solution several times and clean the hole well with a nail or awl.

TIN AND LEAD

Metals are not very convenient for experiments: experiments with them require, as a rule, sophisticated equipment. But some experiments can be done in the home laboratory.

Let's start with tin. Hardware stores sometimes have metal soldering tin sticks. With such a small ingot, you can do an experiment: take a tin stick with both hands and bend it - a distinct crunch will be heard.

Metallic tin has such a crystal structure that when bent, the metal crystals rub against each other, as it were, and a crackling sound occurs. By the way, on this basis, one can distinguish pure tin from tin alloys - an alloy stick does not make any sounds when bent.

And now let's try to get tin from empty cans, from the very ones that are better not to be thrown away, but to be scrapped. Most of the cans are tin-plated on the inside, that is, they are covered with a layer of tin, which protects the iron from oxidation and food from spoilage. This tin can be recovered and reused.

First of all, an empty jar must be properly cleaned. Normal washing is not enough, so pour a concentrated solution of washing soda into the jar and put it on fire for half an hour so that the washing solution boils properly. Drain the solution and rinse the jar two to three times with water. Now you can consider it clean.

We need two or three flashlight batteries connected in series; you can, as mentioned above, take a rectifier with a transformer or a 9-12 V battery. Whatever the current source, attach a tin can to its positive pole (carefully make sure that there is good contact - you can punch a small hole in the top of the can and put a wire in it). Connect the negative pole to some piece of iron, for example, with a large nail cleaned to a shine. Lower the iron electrode into the jar so that it does not touch the bottom and walls. How to hang it - figure it out yourself, this is a simple thing. Pour a solution of alkaline caustic soda (handle with extreme care!) or washing soda into the jar; the first option is better, but requires extreme accuracy in work.

Since the alkali solution will be needed more than once for experiments, we will tell you here how to prepare it. Add washing soda Na 2 CO 3 to the solution of slaked lime Ca(OH) 2 and boil the mixture. As a result of the reaction, caustic soda NaOH and calcium carbonate, i.e. chalk, are formed, practically insoluble in water. This means that in the solution, which, after cooling, must be filtered, only alkali will remain. But back to the tin can experience. Soon, gas bubbles will begin to form on the iron electrode, and the tin from the can will gradually go into solution. But what if it is necessary to obtain not a solution containing tin, but the metal itself? Well, this is also possible. Remove the iron electrode from the solution and replace it with a carbon electrode. Here you will again be helped by an old battery that has served its purpose, in a zinc cup of which there is a network of a carbon rod. Remove it and connect it with a wire to the negative pole of your current source. Spongy tin will settle on the rod during electrolysis, and if the voltage is chosen correctly, this will happen quite quickly. True, it may happen that the tin from one can is not enough. Then take another jar, carefully cut it into pieces with special metal scissors and put it inside the jar in which the electrolyte is poured. Be careful: the cuttings must not touch the carbon rod!

The tin collected on the electrode can be melted down. Turn off the current, take out a carbon rod with sponge tin, put it in a porcelain cup or in a clean metal jar and hold it on fire. Soon the tin will melt into a dense ingot. Do not touch him or the jar until they are cold!

Part of the spongy tin can not be melted down, but left for other experiments. If you dissolve it in hydrochloric acid - in small pieces and with moderate heating - you get a solution of tin chloride. Prepare such a solution with a concentration of about 7% and add, while stirring, a solution of alkali of a slightly higher concentration, about 10%. At first a white precipitate will fall out, but soon it will dissolve in excess alkali. You've got a solution of sodium stannite - the same one you had in the beginning when you started to dissolve the tin from the jar. But if so, then the first part of the experiment - the transfer of metal from the can to the solution - can no longer be repeated, but proceed immediately to its second part, when the metal settles on the electrode. This will save you a lot of time if you want to get more tin from cans.

Lead melts even more easily than tin. Place a few pellets in a small crucible or in a metal can from shoe polish and heat on a flame. When the lead is melted, carefully remove the jar from the heat by grasping the rim with large, secure tweezers or pliers. Pour the lead melt into a plaster or metal mold, or just into a sand hole - this is how you get a homemade lead casting. If, however, the molten lead is further calcined in air, then after a few hours a red coating forms on the surface of the metal - mixed lead oxide; under the name "red lead" it was often used before for the preparation of paints.

Lead, like many other metals, reacts with acids, displacing hydrogen from them. But try to put lead in concentrated hydrochloric acid - it will not dissolve in it. Take another, obviously weaker acid - acetic acid. In it, lead, though slowly, but dissolves!

This paradox is explained by the fact that when interacting with hydrochloric acid, poorly soluble lead chloride PbCl 2 is formed. Covering the surface of the metal, it prevents its further interaction with the acid. But lead acetate Pb (CH 3 COO) 2, which is obtained by reaction with acetic acid, dissolves well and does not interfere with the interaction of acid and metal.

ALUMINUM, CHROME AND NICKEL

With aluminum, we will first set up two simple experiments, for which a broken aluminum spoon is quite suitable. Place a piece of metal in a test tube with any acid, at least with hydrochloric acid. Aluminum will immediately begin to dissolve, vigorously displacing hydrogen from the acid - aluminum salt A1C1 3 is formed. Dip another piece of aluminum into a concentrated solution of alkali, such as caustic soda (carefully!). And again, the metal will begin to dissolve with the release of hydrogen. Only this time another salt is formed, namely the salt of aluminum acid, aluminate NaAlO 2 .

Aluminum oxide and hydroxide simultaneously exhibit both basic and acidic properties, that is, they react with both acids and alkalis. They are called amphoteric. Tin compounds, by the way, are also amphoteric; check it out for yourself if you've already taken the tin out of the can, of course.

There is a rule: the more active the metal, the sooner it oxidizes and corrodes. Sodium, for example, cannot be left in the air at all; it is stored under kerosene. But this fact is also known: aluminum is much more active than, for example, iron, but iron quickly rusts, and aluminum, no matter how much you keep it in air and in water, practically does not change. What is the exception to the rule?

Let's set up an experience. Fix a piece of aluminum wire in an inclined position over the flame of a gas burner or spirit lamp so that the lower part of the wire is heated. At 660 °C this metal melts; one might expect aluminum to start dripping onto the burner. But instead of melting, the heated end of the wire suddenly sags sharply. Take a closer look and you will see a thin case with molten metal inside. This case is made of aluminum oxide Al 2 O 3 , a durable and very heat resistant substance.

The oxide covers the aluminum surface with a thin and dense layer and prevents it from oxidizing further. This property is used in practice. For example, for metal cladding; a thin aluminum layer is applied to the metal surface, aluminum is immediately covered with oxide, which reliably protects the metal from corrosion.

And two more metals with which we will experiment - chromium and nickel. In the periodic table, they stand far apart, but there is a reason to consider them together: metal products are coated with chrome and nickel so that they shine, do not rust. So, the backs of metal beds are usually covered with nickel, car bumpers - with chrome. Is it possible to know exactly what metal the coating is made of?

Let's try to analyze. Peel off a piece of coating from the old part and leave it in the air for a few days so that it has time to become covered with an oxide film, and then place it in a test tube with concentrated hydrochloric acid (handle with care! Acid should not get on hands and clothes!). If it was nickel, then it will immediately begin to dissolve in acid, forming a salt NiCl 2; this will release hydrogen. If the shiny coating is made of chromium, then at first there will be no changes, and only then the metal will begin to dissolve in acid with the formation of chromium chloride CrCl 3. By removing this piece of acid coating with tweezers, rinsing it with water and drying it in the air, the same effect can be observed again after two or three days.

Explanation: a thin oxide film forms on the surface of chromium, which prevents the acid from interacting with the metal. However, it also dissolves in acid, though slowly. In air, chromium is again covered with an oxide film. But nickel does not have such a protective film.

But in this case, why did we keep the metals in air before the first experiment? After all, the chrome was already covered with a layer of oxide! And then, that only the outer side was covered, and the inner side, facing the product, did not come into contact with the oxygen of the air.

EXPERIMENTS WITH COPPER WIRE

Several interesting experiments can be made with copper, so we will devote a special chapter to it.

From a piece of copper wire, make a small spiral and fix it in a wooden holder (you can leave a free end of sufficient length and wind it around a regular pencil). Ignite the spiral in a flame. Its surface will be covered with a black coating of copper oxide CuO. If the blackened wire is immersed in dilute hydrochloric acid, the liquid will turn blue, and the surface of the metal will again become red and shiny. The acid, if it is not heated, does not act on copper, but dissolves its oxide, turning it into a CuCl 2 salt.

But here's the question: if copper oxide is black, why are antique copper and bronze objects covered not with black, but with a green coating, and what kind of coating is this?

Try to find an old copper object, say a candlestick. Scrape some of the green residue off of it and place it in a test tube. Close the neck of the test tube with a cork with a gas outlet tube, the end of which is lowered into lime water (you already know how to prepare it). Heat the contents of the test tube. Drops of water will collect on its walls, and gas bubbles will be released from the gas outlet pipe, from which the lime water becomes cloudy. So it's carbon dioxide. A black powder will remain in the test tube, which, when dissolved in acid, gives a blue solution. This powder, as you can probably guess, is copper oxide.

So, we learned what components the green plaque decomposes into. Its formula is written as follows: CuCO 3 * Cu (OH) 2 (basic copper carbonate). It forms on copper objects, since there is always both carbon dioxide and water vapor in the air. Green plaque is called patina. The same salt is found in nature - it is nothing more than the famous mineral malachite.

We will return to experiments with patina and malachite - in the "Pleasant with useful" section. Now let's turn our attention back to the blackened copper wire. Is it possible to restore its original shine without the help of acid?

Pour pharmacy ammonia into a test tube, heat a copper wire red-hot and lower it into a vial. The spiral will hiss and become red and shiny again. In an instant, a reaction will occur, as a result of which copper, water and nitrogen are formed. If the experiment is repeated several times, then the ammonia in the test tube will turn blue. Simultaneously with this reaction, another so-called complexation reaction takes place - the very complex copper compound is formed, which previously allowed us to accurately determine ammonia by the blue color of the reaction mixture.

By the way, the ability of copper compounds to react with ammonia has been used since very ancient times (since the time when the science of chemistry was not in sight). Ammonia solution, i.e., ammonia, cleaned copper and brass objects to a shine. So, by the way, experienced housewives are doing now; for greater effect, ammonia is mixed with chalk, which mechanically wipes off dirt and adsorbs impurities from the solution.

next experience. Pour a little ammonia into a test tube - ammonium chloride NH 4 Cl, which is used for soldering (do not confuse it with ammonia NH 4 OH, which is an aqueous solution of ammonia). With a red-hot copper spiral, touch the layer of substance covering the bottom of the test tube. Again there will be a hiss, and white smoke will rise up - this is the particles of ammonia escaping, And the spiral will again sparkle with its original copper luster. A reaction took place, as a result of which the same products were formed as in the previous experiment, and in addition copper chloride CuCl 2 .

It is because of this ability - to restore metallic copper from oxide - that ammonia is used for soldering. The soldering iron is usually made of copper, which conducts heat well; when its “sting” oxidizes, copper loses its ability to hold tin solder on its surface. A little ammonia - and the oxide is gone.

And the last experiment with a copper spiral. Pour a little eau de cologne (even better, pure alcohol) into the test tube and add the red-hot copper wire again. In all likelihood, you can already imagine the result of the experiment: the wire has again been cleared of the oxide film. This time, a complex organic reaction took place: the copper was reduced, and the ethyl alcohol contained in the cologne was oxidized to acetaldehyde. This reaction is not used in everyday life, but sometimes it is used in the laboratory when an aldehyde is to be obtained from alcohol.

That's all our first, introductory experiments. Now that you, as they say, have gotten your hands on the experiment, and if you are experimenting at home, then you have probably created a certain supply of dishes and available reagents, it's time to get serious about experiments. Let's take a look at the kitchen cabinet....

Guys, we put our soul into the site. Thanks for that
for discovering this beauty. Thanks for the inspiration and goosebumps.
Join us at Facebook and In contact with

We take care of our children every day - we cook porridge for them in the morning and iron their clothes. But in 20 years, they will remember not our household chores, but the moments spent together.

website collected 16 experiments that will tear adults away from business and captivate children. They do not need a lot of time and some special preparation, and there will be a lot of pleasure. And then you can cook porridge. Together.

solid liquid

You will need:

• starch
• Plastic container
• food coloring, board, hammer and nails for more experiments

Mix water and starch in a container until creamy. You get a "non-Newtonian" fluid. You can easily dip your fingers into it, but if you hit the surface with your fist, you will feel that it is hard. Place a board on the surface of the liquid and you will easily drive in a nail, but it is worth drowning one corner of it in the liquid, and the board will easily sink to the bottom. If desired, the "solid liquid" can be colored with food coloring.

DIY kinetic sand

You will need:

• 4 tsp boric alcohol
• 2 tsp stationery glue
• 1 tsp dye
• 100 g sand for chinchillas
• glass bowl

Pour all liquid ingredients into a bowl, add sand and mix thoroughly. Done, you can create!

pharaoh snake

You will need:

• sand
• alcohol
• sugar
• matches
• plate for "snake"

Pour sand into a plate in a slide, soak it in alcohol, and put a mixture of sugar and soda on top. Set it on fire. "Snake" grows instantly!

Electric train made of wire and batteries

You will need:

• a coil of thick copper wire (the more wire, the longer the "tunnel")
• 1 AA battery
• 2 round neodymium magnets suitable for the battery in diameter
• ordinary pen

Wind the wire around the handle to make a long spring. Attach magnets to both ends of the battery. Start the train. He will drive himself!

Burning candle swing

You will need:

• candle
• thick needle
• lighter
• two glasses
• pliers

Cut off the bottom end of the candle by a centimeter and a half to free the wick. Clamp the needle in pliers and heat it with a lighter, and then pierce the candle in the middle. Put it on the edges of two cups and light it on both sides. Slightly swing, and then the candle will begin to rotate itself.

paper towel rainbow

You will need:

• food colorings
• paper towels
• 5 glasses

Place the cups in a row and pour water into the 1st, 3rd and 5th. In the 1st and 5th drop red food coloring, in the 3rd - yellow, in the 5th - blue. Fold 4 paper towels 4 times to make strips and then fold them in half. Insert the ends into different glasses - one between the 1st and 2nd glass, the second between the 2nd and 3rd, etc. After a couple of hours, you can admire the rainbow!

Elephant Toothpaste

You will need:

• 3/4 cup water
• 1 tsp potassium permanganate
• 1 st. l. liquid soap
• hydrogen peroxide
• glass flask
• disposable gloves

Dissolve potassium permanganate in water, add liquid soap and pour the mixture into a glass flask. Carefully but quickly pour in the peroxide. Stormy foam will splash up from the flask - a real toothpaste for an elephant!

very slow ball

You will need:

• steel ball
• transparent plastic ball-container of two halves
• liquid honey

Put the steel ball into the container, pour in the honey and start the whole structure down the slide. Hmm, what if you try it with shower gel?

smoke rings

You will need:

• plastic bottle (0.5 l)
• balloon
• incense stick
• lighter
• scissors

Cut off the bottom of the plastic bottle and half of the balloon. Put the wide part of the balloon on the cut of the bottle. Insert the wand into the bottle, cover its opening with your hand and wait until it fills with smoke. Start the smoky rings by sharply tapping your finger on the stretched ball.

Self-inflating balloons

You will need:

• 4 plastic bottles
• table vinegar
• 3 art. l. soda
• 3 balloons
• liquid food coloring

Cut off the top of the plastic bottle, pull all the balls one by one onto the hole and pour a spoonful of soda into each ball through the resulting funnel. Pour the vinegar on the bottoms of the bottles, drop the food coloring there and carefully, so that the soda does not spill into the bottle, pull the balls over the holes. It remains to raise them - the soda will pour out, react with vinegar, and the balls will inflate themselves.

Acetic Soda Rocket

You will need:

• plastic bottle (2 l)
• 3 simple pencils
• 2 tbsp. l. soda
• 200 ml vinegar 9%
• wide tape
• wine cork
• paper towel

Make sure that the cork fits snugly on the neck of the bottle. Tape the pencils to the top of the bottle so that it can stand up. Pour the vinegar into the bottle. Wrap the baking soda tightly in a paper towel and twist the ends tightly. Go outside, dip a package of soda into a bottle and plug it with a cork, pressing one end of the package to the neck. Flip the rocket, put it on the ground and run! Takeoff should be observed from 15–20 meters, no less.

The interaction of metals with salts

Active metals displace less active ones from salts (metals are arranged in descending order of activity in a series of voltages).

Let's experiment with a solution of divalent copper sulfate CuSO 4 . In one flask with a solution we put pieces of zinc Zn, in another - steel buttons (steel is an alloy based on iron Fe). What will happen in a few hours? The solutions changed color, which means that there is no more copper sulfate left. Active metals - zinc and iron replaced copper in sulfate and formed salts. Zinc and iron are oxidized, and copper is restored.

CuSO 4 + Zn \u003d ZnSO 4 + Cu

CuSO 4 + Fe \u003d FeSO 4 + Cu

In one flask, copper stood out on the buttons, in the other - on pieces of zinc. There were different metals in the flasks, so the copper deposit looks different. On zinc, copper stood out in the form of a loose brown mass. On iron buttons, the copper deposit is more dense, pink in color.

Equipment: flasks.

Safety. Careful handling of copper salts is necessary. Copper salts in high concentrations are poisonous. Require compliance with the rules for working with toxic substances. Avoid contact with copper salts on the skin and mucous membranes.

Statement of experience– Elena Makhinenko, text– Ph.D. Pavel Bespalov.

The interaction of tin chloride (II) with zinc ("Tin Hedgehog")

More active metals can replace less active metals from solutions of their salts. Pour a solution of tin (II) chloride into a glass, put a zinc plate into the solution. After some time, the plate is covered with a beautiful "fluffy" coating of tin. There was a reduction of tin from a solution of its salt with a more active metal - zinc:

SnCl 2 + Zn = sn + ZnCl 2

Equipment: chemical glass, glass rod.

Safety engineering. The experience is safe.

Statement of experience and text– Ph.D. Pavel Bespalov.

Demonstration of the properties of Wood's alloy.

Wood's alloy consists of four components. It contains 50% bismuth, 25% lead, 12.5% ​​tin and 12.5% ​​cadmium. Drop the alloy granules into hot water. It goes into a liquid state. It is a low melting alloy. The melting point of the alloy is about +70 °C. Meanwhile, the melting point of tin is +232 °C, cadmium +321 °C, bismuth +271 °C, lead +327 °C. The melting point of an alloy differs from the melting points of the metals in its composition.

Equipment: chemical glass, tripod, burner, tweezers.

Safety engineering. Observe the rules for handling heating devices.

Statement of experience and text– Ph.D. Pavel Bespalov.

Platinum - a catalyst for the combustion of hydrogen

At ordinary temperatures, hydrogen very rarely enters into chemical reactions. Hydrogen does not react with oxygen either. But if you direct a stream of hydrogen at finely crushed platinum, then the hydrogen ignites. This property of platinum was used in the so-called "Döbereiner hydrogen flint", which served to produce fire. We will obtain hydrogen in the Kiryushkin apparatus, which is similar in principle to the Kipp apparatus. Let's check the hydrogen for purity. To do this, fill the test tube with escaping hydrogen and bring the test tube to the flame of the burner. A quiet clap indicates the purity of the hydrogen released. With tweezers, take a little platinized asbestos (asbestos coated with finely crushed platinum). Let's direct a jet of hydrogen at platinized asbestos. Asbestos is heated, and hydrogen ignites.

2H 2 + O 2 \u003d 2H 2 O

Equipment: Kiryushkin apparatus, test tube, tweezers, burner.

Safety engineering. Observe the rules for working with combustible gases. Hydrogen can only be used after checking for purity.

Statement of experience and text– Ph.D. Pavel Bespalov.

Self-ignition of nickel in air

Nickel is a strong, corrosion-resistant metal that does not change under the influence of atmospheric oxygen and moisture. Nickel is used to coat the details of devices and machines to give a decorative look and protection against corrosion. But crushed metals, including nickel, differ in their properties from metals that are in a monolithic form. Separate nickel from a nickel-aluminum alloy by placing the alloy powder in an alkali solution.

Aluminum actively reacts with alkali, dissolving in it, the reaction proceeds with the release of hydrogen. To increase the rate of aluminum dissolution, heat the solution. When the reaction is over, and all the aluminum goes into solution, we wash the resulting nickel crumb first with water and then with ethyl alcohol to remove residual moisture. We will extract some nickel crumbs from alcohol on filter paper. When the alcohol evaporates, nickel begins to react with atmospheric oxygen, gradually heats up and burns to form nickel oxide.

2 Ni + O 2 = 2 NiO

Finely divided iron also has similar properties. Crushed nickel and iron are pyrophores. Pyrophores are substances or mixtures of substances that ignite spontaneously in air.

Equipment: chemical beaker, filter paper, grid stand, glass rod.

Safety engineering. Observe the rules for working with alkalis and fire safety rules. All pyrophoric nickel residues should be destroyed by dissolving them in dilute nitric acid.

Statement of experience and text– Ph.D. Pavel Bespalov.

Electrolysis of potassium iodide solution

Electrolysis is the decomposition of a substance under the action of an electric current. The electrolysis of potassium iodide takes place with the release of alkali, hydrogen and iodine:

2KI + 2 H 2 O = 2 KOH + H 2 + I 2

Let's prepare an electrolyzer filled with a solution of potassium iodide, and two test tubes with the same solution. To detect alkali, add a solution of phenolphthalein to one of the test tubes (this test tube is for the cathode), to detect iodine, add starch to another test tube (test tube for the anode). Place the test tubes prepared in this way on the electrodes and turn on the current. In one of the test tubes on the cathode, we observe the evolution of hydrogen, the solution in this test tube becomes raspberry: alkali has formed in the test tube. A blue color appeared in the second test tube. In this test tube, iodine was released as a result of electrolysis. The iodine turned the starch blue. We have seen how during the electrolysis of a solution of potassium iodide, iodine is formed, hydrogen gas and potassium hydroxide are released.

Equipment: test tubes, test tube stand, chemical beakers, pipette, test tube holder, electrolysis device, beaker.

Safety. Follow the rules for working with electrical appliances.

Statement of experience– Elena Makhinenko, text– Ph.D. Pavel Bespalov.

Electrochemical series of voltages - displacement of hydrogen by metals.

Metals differ in chemical activity. Metals are arranged in descending order of activity in a series of voltages:

Li, K, Ca, Na, Mg, Al, Mn, Zn, Fe, Co, Ni, Sn, Pb, H 2 , Cu, Hg, Ag, Au

Active metals (from lithium to lead) reduce hydrogen from acids, inactive metals (from copper to gold) do not.

Let's test four metals: magnesium Mg, aluminum Al, iron Fe and copper Cu. Prepare test tubes with a solution of hydrochloric acid (HCl) and immerse the metals in them. Copper does not react with hydrochloric acid solution. Iron slowly reduces hydrogen from an acid solution. Aluminum reacts more actively with hydrochloric acid solution, reducing hydrogen.

Magnesium most vigorously reduces hydrogen from hydrochloric acid. We have seen that metals standing in the electrochemical series of voltages up to hydrogen (iron, aluminum and magnesium) restore it from acid solutions.

Metals in the row after hydrogen (copper in our experiment) do not reduce it from acids. Magnesium was the most active metal in our experiment, copper was the least active.

2 HCl + Mg \u003d MgC1 2 + H 2

2 HCl + Fe = FeC1 2 + H 2

6 HCl + 2Al = 2 A1C1 3 + 3H 2

Equipment:

Safety. Follow the rules for working with acid solutions. Avoid contact of acids with skin and mucous membranes.

As a result of the reaction, a combustible gas is formed - hydrogen: there should not be an open flame nearby.

Statement of experience– Elena Makhinenko, text– Ph.D. Pavel Bespalov.

Electrochemical series of voltages of metals. Displacement of a metal from salt by other metals

Metals are arranged in descending order of activity in a series of voltages:

Li, K, Ca, Na, Mg, Al, Mn, Zn, Fe, Co, Ni, Sn, Pb, H 2 , Cu, Hg, Ag, Au

Active metals displace less active metals from solutions of their salts. In the first test tube - copper (Cu) and a solution of a salt of a less active metal - silver (AgNO 3). The second pair is iron (Fe) and copper salt solution (CuSO 4). Iron is more active than copper. In the third tube - zinc (Zn) and a salt solution of less active lead - Pb(NO 3) 2 . Reactions begin in the test tubes. After a while, we'll see what happened in the test tubes. Copper covered with white crystals of silver:

2 AgNO 3 + Cu = Cu(NO 3 ) 2 + 2 Ag

A pink coating of metallic copper appeared on the iron nail:

CuSO 4 + Fe = FeSO 4 + Cu

Zinc was covered with a loose layer of metallic lead:

Pb (NO 3) 2 + Zn \u003d Pb + Zn (NO 3) 2

We have seen that active metals displace less active ones from solutions of their salts.

Equipment: test tubes, test tube rack, funnel, tweezers.

Safety. Lead salts and silver salts are poisonous, avoid contact with skin and mucous membranes. Silver nitrate solution leaves black spots on clothing and skin.

Statement of experience– Elena Makhinenko, text– Ph.D. Pavel Bespalov.

Hydrogen peroxide, which is the basis of our experience, is a very unstable compound. A substance consisting of two hydrogen atoms and two oxygen atoms decomposes into oxygen and water even in the absence of any external stimuli. However, this process is very slow. To significantly speed it up, just add a small amount of catalyst. Barely noticeable traces of the presence of copper, iron, manganese, and even ions of these metals can trigger a violent decomposition reaction.

1. Pour 200 ml of a 3% hydrogen peroxide solution into a plastic bottle. Such a solution is sold in a pharmacy as an antiseptic. Instead of peroxide, you can take bleach - they are also prepared on the basis of H2O2.

Hydrogen peroxide (otherwise known as peroxide) is dangerous to living beings. To decompose H2O2 into oxygen and water, an enzyme called catalase is used. Catalase is found in almost all living organisms, including the yeast that we use in our experiments.

2. Add food coloring. It is better to use food-grade paints - not because we are going to eat foam (this is not useful anyway), but because they definitely do not contain catalysts for the decomposition of hydrogen peroxide.

Hydrogen peroxide is a liquid with a density of 1.4 g/cm3. The oxygen released during its decomposition is a gas, one gram of which occupies as much as 700 cm³.

3. Add detergent. Dishwashing detergents are the best. Volume - about half of the volume of peroxide, that is, 100 ml.

Of course, for experiments we use only a 3% solution of hydrogen peroxide, however, this is enough for its decomposition to release gas in a volume much larger than the original one.

4. Dilute the yeast in warm water using a separate glass for this. This is not so easy to do - the yeast will stick together in lumps. You need to patiently stir a tablespoon of yeast into 50 ml of water, and then let it stand for five minutes. Resolutely pour the yeast solution into the hydrogen peroxide bottle and get ready to watch. With luck, the reaction will go so intense that the foam will literally jump out of the bottle.

To see the released oxygen, we catch it in soap bubbles. To do this, add a foaming dishwashing detergent to the hydrogen peroxide solution.