Experiments with copper wire

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

Make a small spiral from a piece of copper wire and secure it in a wooden holder (you can leave a free end of sufficient length and wrap it around a regular pencil). Heat the coil in a flame. Its surface will be covered with a black coating of copper oxide CuO. If a blackened wire is dipped into 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 the salt CuCl 2.

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

Try finding an old copper object, say a candlestick. Scrape off some of the green residue and place it in a test tube. Close the neck of the test tube with a stopper 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 tube, causing the limewater to become cloudy. So it's carbon dioxide. What remains in the test tube is a black powder, which when dissolved in acid gives a blue solution. This powder, as you probably guess, is copper oxide.

So, we found out what components green plaque decomposes into. Its formula is written as follows: CuCO 3 * Cu(OH) 2 (basic copper carbonate). It forms on copper objects because there is always both carbon dioxide and water vapor in the air. The green coating is called patina. The same salt is found in nature - it is none other than the famous mineral malachite.

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

Pour ammonia into a test tube, heat the copper wire red-hot and lower it into the vial. The spiral will hiss and again become red and shiny. In an instant, a reaction will occur resulting in the formation of copper, water and nitrogen. If the experiment is repeated several times, the ammonia in the test tube will turn blue. Simultaneously with this reaction, another, so-called complexation reaction occurs - the same copper complex compound is formed, which previously allowed us to accurately identify 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 (even since those times when the science of chemistry was not even in sight). Copper and brass objects were cleaned with an ammonia solution, i.e., ammonia, to a shine. This, by the way, is what experienced housewives do now; for greater effect, ammonia is mixed with chalk, which mechanically scrubs away dirt and adsorbs contaminants from the solution.

Next experience. Pour a little ammonia-ammonium chloride NH 4 Cl into the test tube, which is used for soldering (do not confuse it with ammonia NH 4 OH, which is an aqueous solution of ammonia). Using a hot copper spiral, touch the layer of substance covering the bottom of the test tube. The hissing will be heard again, and white smoke will rise up - these are the particles of ammonia evaporating, and the spiral will again sparkle with its pristine copper shine. A reaction occurred, as a result of which the same products were formed as in the previous experiment, and in addition copper chloride CuCl 2.

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

And the last experiment with a copper spiral. Pour a little cologne into the test tube (even better - pure alcohol) and again introduce the hot copper wire. In all likelihood, you can already imagine the result of the experiment: the wire was again cleared of the oxide film. This time a complex organic reaction occurred: 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 needs to be obtained from alcohol.

Metals are not very convenient for experiments: experiments with them usually require complex equipment. But some experiments can be carried out in a home laboratory.

Let's start with tin. Hardware stores sometimes sell tin sticks for soldering. You can do an experiment with such a small ingot: take a tin stick with both hands and bend it - there will be a distinct crunch.

U metal tin such a crystalline structure that when bent, the metal crystals seem to rub against each other, producing a crunching sound. By the way, by this feature you can distinguish pure tin from tin alloys - a stick made of an alloy does not make any sounds when bent.

Now let’s try to extract tin from empty tin cans, the very ones that are best not thrown away, but scrapped. Most cans from the inside tinned, i.e. they are coated with a layer of tin, which protects iron from oxidation and food products from spoilage. This tin can be recovered and reused.

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

We will need two or three batteries for a flashlight, 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 ensure that there is good contact - you can punch a small hole in the top of the can and insert a wire into it).

Negative pole connect to some piece of iron, for example, a large nail, cleaned to a shine. Lower the iron electrode into the jar so that it does not touch the bottom and walls. Figure out how to hang it yourself, it’s a simple thing. Pour the lye solution into the jar - caustic soda (handle with extreme caution!) or washing soda; The first option is better, but requires extreme care in work.

Since an alkali solution will be needed for experiments more than once, we will tell you here how to prepare it. Add washing soda Na 2 CO 3 to a solution of slaked lime Ca(OH) 2 and boil the mixture. As a result of the reaction, caustic soda NaOH and calcium carbonate are formed, i.e. chalk, practically insoluble in water. This means that in the solution, which after cooling must be filtered, only alkali will remain. But let's return to the experience with the tin can. Soon on the iron electrode Gas bubbles will begin to appear, and the tin from the tin can will gradually become go into solution.

Well, what if you need to get not a solution containing tin, but the metal itself? Well, this is possible too. Remove the iron electrode from the solution and replace it with a carbon one. Here, an old, worn-out battery will help you again, with a carbon rod in its zinc cup. Remove it and connect the wire to the negative pole of your power source. Spongy tin will settle on the rod during electrolysis, and if the voltage is selected 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 place it inside the jar into which the electrolyte is poured. Be careful: the cuttings must not touch the carbon rod!

Assembled on the electrode tin can be melted down. Turn off the current, take out a charcoal rod with sponge tin, put it in a porcelain cup or a clean metal can and hold it on the fire. Soon the tin will be fused into a dense ingot. Do not touch it or the jar until they have cooled!

Part of the sponge 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 will get a solution tin chloride. Prepare such a solution with a concentration of approximately 7% and add, stirring, an alkali solution of a slightly higher concentration, about 10%. At first a white precipitate will form, but it will soon dissolve in the excess alkali. You got the solution sodium hydroxostannate- the same one that formed in the beginning when you began to dissolve tin from a jar.

But if so, then the first part of the experiment - transferring the metal from the jar into the solution - can no longer be repeated, but proceed immediately to the 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 metal shoe polish can and heat over a flame. Once the lead has melted, carefully remove the jar from the heat by grasping the side of the jar with large, secure tweezers or pliers. Pour the melted lead into a plaster or metal mold, or simply into a sand hole - this way you will get a homemade lead casting. If you continue to calcinate molten lead in air, then after a few hours a red coating will form on the surface of the metal - double lead oxide; entitled " red lead"It was often used in the past to make paints.

Lead, like many other metals, interacts with acids, displacing hydrogen. But try putting lead in concentrated hydrochloric acid- he will not dissolve in it. Take another, obviously weaker acid - vinegar. Lead in it, although slowly, dissolves!

This paradox is explained by the fact that when interacting with hydrochloric acid, a poorly soluble lead chloride PbCl2. By covering the surface of the metal, it prevents its further interaction with the acid. And here 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.

WITH aluminum We will first carry out 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 salt. 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 alkali solution, for example caustic soda(carefully! ). And again the metal will begin to dissolve with the release of hydrogen. Only this time another salt is formed, namely: sodium aluminate.

Oxide And aluminum hydroxide exhibit both basic and acidic properties at the same time, i.e. they react with both acids and alkalis. They are called amphoteric. Connections tin, by the way, are also amphoteric; check it out for yourself, assuming you've already removed the tin from the tin.

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

Let's put experience. Secure the piece aluminum wire in an inclined position over the flame of a gas burner or alcohol lamp so that the lower part of the wire is heated. At 660 o C this metal melts; it would seem that one might expect that aluminum will begin to drip 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 containing molten metal. This "case" is from aluminum oxide Al 2 O 3, a durable and very heat-resistant substance.

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

And two more metals with which we will experiment - chromium And nickel. In the periodic table they are far apart from each other, but there is a reason to consider them together: metal products are coated with chromium and nickel so that they shine and do not rust. Thus, the backs of metal beds are usually covered with nickel, car bumpers - with chrome.

Is it possible to know for sure what metal is the coating made of?? Let's try to analyze. Break off a piece of the coating from the old part and leave it in the air for several days so that it has time to become covered with a film of oxide, and then place it in a test tube with concentrated hydrochloric acid (handle with care! Acid should not get on your hands or clothes!).

If it was nickel, then it will immediately begin to dissolve in the acid, forming the 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 the acid with the formation chromium chloride CrCl 3. By removing this piece of coating from the acid with tweezers, rinsing it with water and air drying it, after two or three days you can observe the same effect again.

Explanation: a thin oxide film forms on the chromium surface, which prevents the interaction of acid with metal. However, it also dissolves in acid, albeit 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 the air before the first experiment? After all, the chromium was already covered with a layer of oxide! And then, only the outer side was covered, and the inner side, facing the product, did not come into contact with oxygen in the air.

With copper you can put several curious experiments, so we will devote a special chapter to it.

Make a piece of copper wire small spiral and secure it in a wooden holder (you can leave a free end of sufficient length and wrap it around a regular pencil). Ignite spiral in flame. Its surface will be covered with a black coating copper oxideСuO. If a blackened wire is dipped into diluted hydrochloric acid, the liquid will turn blue, and the metal surface 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 the salt CuCl 2.

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

Try to find old copper object, say, a candlestick. Scrape off some of the green residue and place it in a test tube. Close the neck of the test tube with a stopper with a gas outlet tube, the end of which is lowered into lime water(you already know how to cook 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 tube, from which lime water becomes cloudy. So this is carbon dioxide. What remains in the test tube is a black powder, which when dissolved in acid gives a blue solution. This powder, as you probably guess, copper oxide.

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

Let's pay attention to blackened copper wire. Is it possible to return it to its original shine without the help of acid? Pour the pharmacy into the test tube ammonia, heat the copper wire red-hot and lower it into the vial. The spiral will hiss and again become red and shiny. In an instant, a reaction will occur resulting in the formation of copper, water And nitrogen. If the experiment is repeated several times, the ammonia in the test tube will turn blue. Simultaneously with this reaction, another reaction occurs, the so-called complexation reaction- the same copper complex compound is formed, which previously allowed us to accurately identify 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 (even since those times when the science of chemistry was not even in sight). Copper and brass objects were cleaned with an ammonia solution, i.e., ammonia, to a shine. This, by the way, is what experienced housewives do now; for greater effect, ammonia is mixed with chalk, which mechanically removes dirt and adsorbs contamination from solution.

Next experience. Pour some into the test tube ammonia - ammonium chloride NH 4 Cl, which is used for soldering (do not confuse it with ammonia, which is an aqueous solution of ammonia). Using a hot copper spiral, touch the layer of substance covering the bottom of the test tube. There will be a hissing sound again, and white smoke will rise up - this is the ammonia particles evaporating. And the spiral will again sparkle with its pristine copper shine. A reaction occurred, as a result of which the same products were formed as in the previous experiment, and in addition copper chlorideСuСl 2.

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

AND last experience with a copper spiral. Pour some cologne into the test tube (even better - pure alcohol) and reintroduce the hot copper wire. In all likelihood, you can already imagine the result of the experiment: the wire was again cleared of the oxide film. This time it happened complex organic reaction: copper recovered, and ethanol contained in cologne has oxidized to acetaldehyde. This reaction is not used in everyday life, but sometimes it is used in the laboratory when an aldehyde needs to be obtained from alcohol.

You can perform several interesting experiments with copper, so we will devote a special page to it.

Make a small spiral from a piece of copper wire and secure it in a wooden holder (you can leave a free end of sufficient length and wrap it around a regular pencil). Heat the coil in a flame. Its surface will be covered with a black coating of copper oxide CuO. If a blackened wire is dipped into 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 the salt CuCl 2.

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

Try finding an old copper object, say a candlestick. Scrape off some of the green residue and place it in a test tube. Close the neck of the test tube with a stopper 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 tube, causing the limewater to become cloudy. So it's carbon dioxide. What remains in the test tube is a black powder, which when dissolved in acid gives a blue solution. This powder, as you probably guess, is copper oxide.

So, we found out what components green plaque decomposes into. Its formula is written as follows: CuCO 3 * Cu (OH) 2 (basic copper carbonate). It forms on copper objects because 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 .

Let's turn our attention back to the blackened copper wire. Is it possible to return it to its original shine without the help of acid?

Pour ammonia into a test tube, heat the copper wire red-hot and lower it into the vial. The spiral will hiss and again become red and shiny. In an instant, a reaction will occur resulting in the formation of copper, water and nitrogen. If the experiment is repeated several times, the ammonia in the test tube will turn blue. Simultaneously with this reaction, another, so-called complexation reaction occurs - the same copper complex compound is formed, which previously allowed us to accurately identify 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 (even since those times when the science of chemistry was not even in sight). Copper and brass objects were cleaned with an ammonia solution, i.e., ammonia, to a shine. This, by the way, is what experienced housewives do now; for greater effect, ammonia is mixed with chalk, which mechanically scrubs away dirt and adsorbs contaminants from the solution.

Next experience. Pour a little ammonia-ammonium chloride NH 4 Cl into the test tube, which is used for soldering (do not confuse it with ammonia NH 4 OH, which is an aqueous solution of ammonia). Using a hot copper spiral, touch the layer of substance covering the bottom of the test tube. There will be a hissing sound again, and white smoke will rise up - this is the ammonia particles evaporating. And the spiral will again sparkle with its pristine copper shine. A reaction occurred, as a result of which the same products were formed as in the previous experiment, and in addition copper chloride CuCl 2.

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

And the last experiment with a copper spiral. Pour a little cologne into the test tube (even better - pure alcohol) and again introduce the hot copper wire. In all likelihood, you can already imagine the result of the experiment: the wire was again cleared of the oxide film. This time a complex organic reaction occurred: 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 needs to be obtained from alcohol.

O. Holguin. "Experiments without explosions"
M., "Chemistry", 1986

What is especially good about this experience is that you probably have everything you need for it at home: a candle, a pharmaceutical potion (alcohol solution, iodine tincture) and some worthless iron object - an old door hinge, a key to an unknown lock or a lock whose keys have been lost . Sand the metal surface on which the design will be made with sandpaper until it shines, light a candle and tilt it so that the paraffin drips onto the shiny surface. Heat the object slightly, then the paraffin will spread into a thin layer. And when it cools down and cools down, use a needle to scratch the grooves so that they reach the metal. Pipette some pharmaceutical iodine and drop it onto the scratches. After a few minutes, the iodine solution will turn pale, and then you need to apply it again to the scratches. After about an hour, remove the layer of paraffin: you will see clear marks on the metal, they exactly repeat the pattern on the paraffin.

If the experience was successful, you can move on to a more serious activity - not just scratch the paraffin, but write a word on it or make a drawing, for example, mark your pocket knife 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 can be 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 it is usually not iodine that 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 etch various objects made of copper and copper alloys, for example, brass. You can try it.

HOMEMADE 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 environment, 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 prepare them ourselves. The starting materials will be plants: many flowers, fruits, berries, leaves and roots contain colored substances that can change their color in response to one or another influence. And when they find themselves in an acidic (or, conversely, alkaline) environment, they visually signal us about this.

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; pick 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 naturally quickly deteriorate - they turn sour and mold. They must be prepared immediately before the experiment. Take some of the stored raw material (the exact amount does not matter), put it in a test tube, add water, place it in a water bath and heat until the solution turns color. After cooling, filter each solution and pour into a clean bottle with a label prepared in advance.

To provide yourself with indicators for the whole year, dry the petals and berries in the summer, place 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 and 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, and soda - green-blue.

Carefully record the results of all these experiments, 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 respond to changes in acidity by changing color. Ordinary borscht can serve as an indicator. Housewives have long noticed this and use this property of beet broth, but not for analysis. To make the borscht bright red, add a little food acid - acetic or citric - to it before the end of cooking; the color changes literally before our eyes.

The indicator phenolphthalein is widely used in laboratories. Let's prepare it from pharmaceutical tablets of the same name. Grind one or two tablets and dissolve in about 10 ml of vodka (in extreme cases, just 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 be useful, and more than once, for determining an alkaline environment: in it it instantly turns red. To check, add a drop or two of phenolphthalein to the washing soda solution.

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

We invite you to continue the table yourself.

And one last thing about plant indicators. It was once fashionable to write invitations on flower petals; and they were written, depending on the flower and the desired color of the inscription, with a solution of acid or alkali, using a thin pen or a pointed stick. Try, if you want, to write 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 called extraction.

Grind a few nut kernels and a handful of sunflower seeds (without husks, of course), put them in a test tube and fill them with gasoline. There should be no fire nearby - gasoline may catch fire! Shake the test tube and let it stand for two hours, remembering to shake from time to time. Then drain the solution on a saucer and expose it to a draft. As 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 making oil from other seeds. Just don’t even try to taste it!

Another experiment - with leaves. For this 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 a 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 place a glass with a leaf inside. Some time later, remove the leaf with tweezers: it has become discolored and the alcohol has become emerald in color. This is how you extracted chlorophyll - the green pigment of 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 a cream or sauce. This is what they do in food factories: green edible dye is extracted from leaves. To speed up this process, we advise you to first chop the leaves and shake the vessel from time to time.

Another experience. Pour approximately 1 ml of pharmaceutical tincture of iodine into a test tube half filled with water; the result will be a brownish solution. Add an equal amount of gasoline to it, shake it several times and leave it alone. When the mixture stratifies, it turns out that the upper, gasoline layer has become dark brown, and the lower, water layer is almost colorless. Iodine dissolves poorly in water, but well in gasoline. That's why it went from an aqueous solution to a gasoline solution.

Our latest extraction experience is based on the difference in solubility. How to quickly distinguish coffee powder from chicory powder? By smell, this is understandable, but what if the smell is weak or you don’t 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 oka will remain almost colorless. On the contrary, coffee substances easily dissolve in water, and its powder slowly sinks to the bottom, leaving a brown mark behind it.

EXPERIMENTS WITH GASES

We've already worked a little with liquids, let's move on to gases. This is somewhat more difficult, and first of all we will need plugs with holes and gas outlet tubes.

The tube can be glass, metal or even plastic. It is better not to take a rubber stopper - it is difficult to drill holes in it. Take cork or polyethylene plugs - holes in them can be burned with a heated awl. Insert a tube into this hole - for example, from an eye dropper; it should fit into the hole of the plug tightly, without gaps, so the hole in the plug must first be made a little smaller than required, and then gradually expand it, adjusting it to the diameter of the tube. Place a rubber or polyethylene flexible tube 30 centimeters long onto the glass tube, and also insert a short glass tube into the other end.

Now the first experiment 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 it sit. The transparent precipitate above the settled solution is limewater. Carefully drain the sediment; This laboratory technique, as you remember, is called decanting.

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

Take a chilled bottle of mineral water or lemonade. Open the stopper, quickly insert the stopper 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 be released from it. This is carbon dioxide CO 2 (also known as carbon dioxide, carbon dioxide). It is added to water to make it tastier.

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

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 one that causes limewater to become 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 the stopper and tube, and continue to pass carbon dioxide through the limewater. Some time later, the solution will again become transparent, because carbon dioxide reacts with calcium carbonate, turning it into another salt - Ca(HCO 3) 2 bicarbonate, and this salt is very soluble in water.

The next gas we'll look at is one that was recently mentioned: ammonia. It is easily recognized by its sharp characteristic odor - the smell of pharmaceutical 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 the other end. Warm the bottle in warm water. Ammonia vapor is lighter than air and will soon fill the inverted test tube. Still holding the test tube upside down, carefully lower it into the glass of water. Almost immediately, the water will begin to rise up into the test tube, because ammonia dissolves well in water, freeing up space for it in the test tube.

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

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

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 soda rather than ammonia into the bottle. Dip a small, pea-sized piece of calcium carbide carefully wrapped in blotting paper into it and insert the plug with the tube. When the blotting paper gets wet, gas will begin to be released, which you will collect in an inverted test tube as before. A minute later, turn the test tube upside down and hold a lit match. The gas will flare up and burn with a smoky flame. This is the same acetylene that gas welders use.

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

The following experiment with gases can be carried out only with good ventilation, and if there is none, then in the fresh air. We will receive sharp-smelling sulfur dioxide (sulfur dioxide) SO 2.

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

If you are unable to buy sodium sulfite, replace it with the contents of a large cartridge of regular photo developer. True, in this case there will be an admixture of carbon dioxide in the sulfur dioxide, but this will not interfere with the experiment.

OXIDATION-REDUCTION

The experiment with sulfur dioxide showed us one of the many redox reactions. In such reactions, atoms of some substances gain electrons, while others give up electrons. The former are called oxidizing agents (potassium permanganate), the latter are called reducing agents (sulfur dioxide).

Let's do a few more experiments with oxidation - reduction.

Drop diluted iodine tincture on a fresh cut of potato: a blue color will appear. It is the starch found in potatoes that turns blue in the presence of free iodine. This reaction is often used to detect starch, which means it is also a qualitative reaction.

Pour a little sodium sulfite solution onto the same place where you dropped the iodine tincture. The color will fade quickly. This is what happened: the 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's another interesting experiment 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 test tube. The contents of the first test tube will become almost colorless, the second - brownish. In the third test tube, brown flakes will fall out, and in the fourth too, but there will be much more sediment. 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, throw a crystal of potassium permanganate into an alkali solution - into a concentrated boiled solution of washing soda - and instead of the usual pink color, green will appear.

This experiment turns out even more beautiful when they work 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 heat up, add a fairly concentrated solution of sodium hydroxide. Observe the color of the liquid - first it will become increasingly purple, then blue as the alkalinity increases, 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 experience 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 swamp. Add a little solution of an oxidizing agent - potassium permanganate - to the test tubes. In tap water it will remain pink, in water from a puddle it will become discolored. In warm weather, organic matter accumulates in standing water. They, like sodium sulfite, reduce potassium permanganate and 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, 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 a little bleach (this is a common disinfectant and must be handled with care). The contents of the test tube will turn yellow. Chloride of lime is a good oxidizing agent; it oxidizes hydroquinone to quinone, which is colored yellow. If you now add a mixture of sodium sulfite and soda from a large cartridge to the test tube, the yellow color will disappear: sodium sulfite will again reduce the quinone to hydroquinone.

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

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

So, you threw a little zinc into a glass with a solution, and the dichromate, being reduced, changes color to dark green. This resulted in the formation of Cr 3+ ions. At the same time, due to the reaction of zinc with acid, a 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; As you do this, oxidation will occur and the solution will turn green again.

ADSORPTION

Probably everyone is familiar with the physicochemical phenomenon that we will now discuss, although perhaps not everyone knows that it is called adsorption. Even if you did not study adsorption in class, you have observed it more than once. As soon as you put an ink blot on paper or, what is much worse, on clothing, you immediately become familiar with this phenomenon. When the surface of one substance (paper, fabric, etc.) absorbs particles of another substance (ink, etc.), this is adsorption.

A very good adsorbent is coal. And not stone, but wood, and not just wood, but active (activated). This type of charcoal is sold in pharmacies, usually in tablet form. This is where we will begin our adsorption experiments.

Prepare a pale solution of ink of any color and pour it into the test tube, but not to the top. Place a tablet of active carbon, preferably crushed, into the test tube, close it with your finger and shake well. The solution will lighten 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 less easily.

There is nothing strange about this: activated carbon differs from regular carbon in that it has a much larger surface area. Its particles are literally riddled with pores (for this purpose, coal is processed in a special way and impurities are removed from it). And since adsorption is absorption by a surface, it is clear: the larger the surface, the better the absorption. Adsorbents are capable of absorbing substances not only from solutions. Take a half-liter glass jar and add one drop of cologne or any other odorous substance to the bottom. Place your palms around the jar and hold it there for half a minute to warm up 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 light waves of your hand, direct the air along with the vapor of the substance to your nose; It’s not always known whether the substance in the bottle smells good.

Whatever the smell, you will, of course, feel it clearly. Now put some active carbon in the flask, close it tightly with a lid and leave for a few minutes. Remove the lid and again direct the air towards you with palm waves. The smell disappeared. 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 activated 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 foreign odors. It is not for nothing that it is not recommended to keep wheat bread in the same package as rye bread - their smells mix, and each loses its special, unique aroma.

A very good adsorbent is puffed corn, or corn sticks, so beloved by many of us. Of course, there is no point in spending a package or even a quarter of a package on experience, but a few grand... 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 you can no longer eat chopsticks.

Let's return to the experiment with the production of carbon dioxide (carbon dioxide). Fill two test tubes with this gas, place corn sticks in one and shake several times. Next, as before, do the experiment with lime water (you can simply “pour” gas into it from test tubes - 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, not treated with an adsorbent, was “poured.” And from the other test tube, the one with corn sticks, carbon dioxide cannot be removed: it was absorbed by the adsorbent.

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

Nowadays, in many kitchens, various devices are installed above gas stoves to purify the air from fumes and smoke. In such devices, among other things, there is a cartridge with some kind of adsorbent through which contaminated air is driven. What happens in this case, 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 when dry cleaning and removing stains, they most often use exactly the same processes that you recently became acquainted with in experiments. Namely: extraction, oxidation - reduction and adsorption.

Of course, you shouldn’t get your clothes dirty for the sake of experiments. Let's do this: we'll prepare several pieces of light-colored fabric, put different stains on it and try to remove them. And if the experiments are successful, you can take the risk of cleaning your own suit (or someone else’s - if allowed...).

The most common stains are grease stains. They are removed, as a rule, by extraction, selecting a suitable solvent for this. Gasoline, turpentine, and medicinal ether are suitable for removing fresh grease stains. Use a cotton swab soaked in solvent to wipe the stain several times, and the grease will dissolve into the solution. To prevent a halo from remaining on the fabric, it must be wiped with soapy water or a solution of washing powder.

Old grease stains are more difficult to remove; solvent alone is not enough; mixtures are needed. For example, gasoline, medicinal ether and turpentine (7:1:2) or wine alcohol, turpentine and medicinal ether (10:2:1).

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

Oil varnish stains are easily removed by a paste of gasoline and white clay. The dough-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 and absorbs substances extracted by gasoline.

First moisten a fresh oil paint stain with turpentine (to soften it), and then remove it with gasoline. If such treatment can damage the paint, then wipe the stain with a hot solution of glycerin or its mixture with an equal amount of wine alcohol.

Extraction can also remove grass stains. Remember the experiment in which we extracted chlorophyll with alcohol? So, if you wipe the stained area with alcohol (or medicinal ether), you can gradually extract chlorophyll from the stain, and it will become discolored.

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

And in this case we combined extraction with adsorption. In general, when removing stains, such a double technique often turns out to be the most effective: white clay, chalk and similar powders do not allow the tinted solution to spread across 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 have an effect, then these stains on white fabrics can be bleached with a solution of hydrogen peroxide (you can dissolve a hydroperite tablet in half a glass of water). Soak the stain with this solution, adding a few drops of ammonia to it, wipe with a clean cotton swab and rinse with water. Hydrogen peroxide (peroxide) is a strong oxidizing agent; it oxidizes many dyes, and they become discolored.

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 fabric overheats, brown thermal oxidation products are formed, and bleach destroys them and makes 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 soda solution to neutralize the acid, and then rinse with clean water.

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

From dry cleaning it would be quite natural to move on to washing, which is what we will do.

Washing is a physical and chemical process, its main characters are surfactants. The molecules of such substances consist of two parts - hydrophilic, i.e., having an affinity for water, and hydrophobic, which does not interact with water, but readily comes into contact with pollutants, for example, difficult-to-clean fats and oils. These groups - hydrophilic and hydrophobic - are located at different ends of a long molecule. Such molecules are attached with their hydrophobic ends to the fatty surface, and the hydrophilic ones stick out, like the needles of a hedgehog. Water wets these “needles” well, it surrounds such a “hedgehog”, tears it off the surface and carries it away. Soap and washing powder work in much the same way. And in order to quickly remove dirt from 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 and add the phenolphthalein solution to the test tube. The color will turn crimson-red. This means 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, breaking down into acid and alkali. But fatty acids are weak, and alkalis in this case are strong, so the solution has an alkaline reaction.

Previously, it was thought that soap washed and washed well because it formed an alkali. It turned out that this was not the case at all. In contrast, alkali (such as washing soda) cleanses because it combines with fats and forms soap-like surfactants in the solution.

By the way, soap is not so difficult 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 stops dissolving. Instead of oil, you can use beeswax. Add a pinch of table salt to the resulting solution. Soap factories do the same thing - this process is called salting out. After adding salt, the solid soap floats to the surface and is easy to separate from the solution.

Nowadays, soap is used for washing less and less, and washing powders are used more and more often. These powders contain surfactants obtained synthetically. That's why they are called synthetic detergents.

Let's do this experiment. Cut a piece of dirty cloth into three parts and place each piece into glasses. Pour just heated water into the first glass, a soap solution into the second, and a solution of any washing powder you can find at home into the third. Lightly rub the scraps, rinse them in clean water, dry them and examine them carefully. That piece of fabric that had been in water did not become much cleaner. The patch of soap solution became noticeably lighter. But the cleanest piece of fabric will be the one you removed from the glass with the washing powder solution. This means that synthetic detergents are more powerful than regular soap.

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

Take plain 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. 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.

Take a piece of dirty cloth again and try washing it with soap and hard water. Nothing will work out for you - not even foam will form. Hardness salts react with soap, calcium and magnesium soaps are formed, and they are insoluble in water. And our soap loses all its beneficial properties.

But if you dissolve washing powder, for example “Lotus”, in hard water, it will wash away dirt almost the same way as before - hard water does not harm it. The surfactants included in the powder do not interact with hardness salts, which means they do not lose their properties.

Solutions of washing powders, like 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 products; 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 you have doubts about whether it is worth washing a woolen sweater with this or that powder, then test with phenolphthalein. The solution has turned red, which means it contains free alkali, which is contraindicated for wool because it can destroy the fibers. But if the solution remains colorless or only slightly colored, 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 away dirt. Try it yourself to see how these substances work. You can take mustard powder or bean decoction 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 detergent effect (you may have come across this expression in old books - soap root). All these natural substances erase, of course, worse than soap, but you can easily see that they still erase.

We will end 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.

Make a flat spiral of several turns from thin copper wire, lightly lubricate it with oil or Vaseline and very carefully lower it to the surface of the water. The surface tension of the water prevents the spiral from sinking, and the water does not wet it. Now, using a pipette, carefully drop one drop of soap solution into the very middle of the spiral. The spiral will immediately begin to spin. Spreading over the surface, the soap solution reaches the end of the spiral, exits it and develops a small jet thrust. When the spiral stops, drop the soap solution again and 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, 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 to 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”, and the “head” tends to water, and the “tail”, on the contrary, repels from the water...

Let's take a closer look at the hydrophobic “tail” - a long hydrocarbon chain. These types of connections are very common and extremely important for industry. They are an indispensable component of many fats, oils, lubricants and other beneficial substances. We will now obtain one of them - the so-called stearin, using laundry soap as a basis.

Using a knife, cut out half a piece of laundry soap and place it in a clean tin can (or in a used saucepan). Pour enough water to cover the soap shavings and place the mixture in a water bath. Stir the contents of the saucepan from time to time with a wooden stick so that the soap dissolves in the water as quickly as possible. When this finally happens, remove the vessel from the heat (not with your bare hand, of course) and pour the vinegar into it. Under the action of acid, a thick white mass will separate 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 making the soap.

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

When the jar has cooled completely, scoop the stearin from the surface with a spoon and transfer it to a clean container. Rinse the stearin two or three times with water and wrap it in a clean white rag or filter paper to absorb excess moisture. When the stearin is completely dry, let's start making the candle.

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

There is a simpler way: immediately coat the wick with stearin heated until softened (you can even just prepare it, not yet cooled down). But in this case, the wick will be less saturated with the fusible mass and the candle will not turn out very good, although it will burn.

For beautiful, shaped candles, the manufacturing methods are not easy. And first of all, you need to make a mold - wooden, plaster, metal. In this case, it is advisable to first soak the wick with one or two layers of stearin; it is then secured in the mold so that it runs exactly down the middle. It is advisable that the wick be slightly stretched. And after that, hot stearin is poured into the mold.

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

Having received a candle from soap, we will conduct the experiment in the opposite direction: we will prepare soap from a candle. But not from paraffin soap; 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 make laundry soap from it. Natural beeswax is also suitable.

Heat several pieces of 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. Keep it in the water bath for a few more minutes, and then, putting on a mitten or wrapping your hand in a towel so as not to get burned, pour the still hot mass into some form - at least into a matchbox. When the soap has hardened, remove it from the box.

Making sure that it is soap and that it cleans is not difficult. Just please don’t use it to wash your hands - we don’t know how pure the substances that made up the candle were.

CHALK, MARBLE, SHELLS...

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

This is a well-known phenomenon: calcium, which is part of chalk, makes the flame red. But why acid? Reacting with chalk, it 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 does not work - it contains an admixture of soda (sodium salt), and the flame turns orange. The best experience is obtained with a piece of white marble soaked in the same acid. And you can make sure that sodium salts color the flame an intense yellow color by adding a grain of NaCl salt to the flame (or simply lightly “salting” the fire).

For the next experiment with chalk, you will need a candle. Strengthen it on a non-flammable stand and add a piece of chalk (marble, shell, eggshell) to the flame. The chalk becomes covered with soot, which means the flame temperature 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 medicine pipette and replace it with a rubber or plastic tube. Blow into the tube so that air enters the flame just above the wick through the drawn end of the pipette. The flame will deviate to the side, its temperature will increase. Point the tongue at the sharpest part of the crayon. This area will become white hot, the chalk here 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, and eggshells. Place the burned pieces in a clean tin. While they are cooling, place the largest piece in a saucer and drop some water on the place that was heated. There will be a hissing sound, all the water will be absorbed, and the baked area will crumble into powder. This powder is slaked lime Ca(OH)2.

Add more water and drop in the phenolphthalein solution. The water in the saucer will turn red; This means that slaked lime forms an alkaline solution.

When the burnt 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 get lime water. Let the liquid settle and pour the clear solution into a clean bottle. Pour some limewater into a test tube - and you can use it to perform the previously described experiments with gases. Or 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

You will encounter experiments with electricity more than once in this book. Now - the simplest ones. To carry them out, three or four flashlight batteries are enough.

In fact, they often try to carry out experiments in electrochemistry at home, but they don’t always work out: some little thing and nothing happens. If you follow all our instructions, you can be sure that the experiment will be a success.

Let's start with a very simple, but nevertheless instructive experience. It requires only one reagent: ink of any color. True, you will 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 fit freely into a transparent vessel - a tall beaker or a large test tube. Before the experiment, drill holes in the plates on one side to attach the conductors. Prepare two identical, literally a few millimeters thick, plastic or wooden spacers and glue them with metal strips so that they are parallel and 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). Place a structure 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 become lighter, and dark particles will collect at the bottom and top.

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

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

In the next experiment, a thick-walled tea glass, expanding at the top, will serve as an electrolytic bath. Prepare a plywood circle of such a diameter that it presses against the wall of the glass three to four centimeters above the bottom. Drill two holes in the mug in advance (or cut a slot in diameter in it), and pierce two holes nearby with an awl: the wiring will pass through them. Insert two pencils 5–6 cm long, sharpened at one end, into the large holes or slot. Pencils, or rather their leads, will serve as electrodes. Make nicks on the raw ends of the pencils to expose the leads, and tape the exposed ends of the wires to them. Twist the wires and carefully wrap them with insulating tape, and for the insulation to be completely reliable, it is best to hide the wires in rubber tubes. All the parts of the device are ready, all that remains is to assemble it, that is, insert the circle with the electrodes inside the glass.

Place the glass on a plate and pour a solution of washing soda ash Na 2 CO 3 into it to the brim 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. Underwater, place the test tube on the pencil electrode. Do the same with the second test tube.

The batteries - at least three in number - must be connected in series, the “plus” of one to the “minus” of the other, and the wires from the pencils must be connected to the outer batteries. Electrolysis of the solution will begin 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 the pencil connected to the minus side has a full test tube of hydrogen, you can take it out and, without turning it over, ignite the gas. It will light up with a characteristic sound. Oxygen is released at the other electrode, the positive one (anode). Close the test tube filled with it with your finger under water, remove it from the glass, turn it over and introduce a smoldering splinter - it will light up.

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

With the same device, you can perform another experiment - electrolysis of a saturated solution of sodium chloride NaCl. In this case, one test tube will be filled with colorless hydrogen, and the other with a 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.

Cover the test tube with chlorine, which also contains a small amount of salt solution, with your finger under water, turn it over and shake without removing your finger. A chlorine solution is formed in a test tube - chlorine water. It has strong whitening properties. For example, if you add chlorine water to a pale blue ink solution, it will become discolored.

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

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

Using current and a saturated solution of table salt, you can perform another entertaining experiment. Let's now start drilling metal with an ordinary pencil.

Prepare a saturated solution of table salt in a tea saucer. Connect the safety razor blade with a wire to the positive terminal of the flashlight battery (the blade will be the anode). Break off the lead at the sharpened end of the pencil and dig it out about half a millimeter with a needle. 2–3 cm higher, make a notch with a knife up to the stylus and wrap the end of the bare wire around it; Wrap this place with insulating tape, and connect 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 with the 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 after ten to fifteen minutes there will be a through hole 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 literally in 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 coat the workpiece with varnish in advance, and remove the varnish where you will drill.

The recess in the lead was needed so that the lead would not touch the metal. Otherwise, the circuit will close immediately, the current will not flow 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). Place the anode plate on a board or plate, drop a drop of water, dip a pencil attached to the battery in salt and immerse its sharpened end in the drop. From time to time, remove electrolysis products with a cloth and apply a new drop. By repeating this operation, you can drill through metal foil or tin from a tin can without any effort. By the way, you can also make a hole in a broken steel knife in order to attach a new handle to it.

Of course, for drilling metal more than a millimeter thick, 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 electrolysis method, you will have to change the electrolyte solution several times and clean the well well with a nail or an awl.

TIN AND LEAD

Metals are not very convenient for experiments: experiments with them usually require complex equipment. But some experiments can be carried out in a home laboratory.

Let's start with tin. Hardware stores sometimes sell tin sticks for soldering. You can do an experiment with such a small ingot: take a tin stick with both hands and bend it - you will hear a distinct crunch.

Metal tin has such a crystalline structure that when bent, the metal crystals seem to rub against each other, producing a crunching sound. By the way, by this feature you can distinguish pure tin from tin alloys - a stick made of an alloy does not make any sounds when bent.

Now let’s try to extract tin from empty tin cans, the very ones that are best not thrown away, but scrapped. Most cans are tinned on the inside, i.e. they are coated with a layer of tin, which protects the iron from oxidation and food products from spoilage. This tin can be recovered and reused.

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

We will 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 ensure that there is good contact - you can punch a small hole in the top of the can and insert a wire into it). Connect the negative pole to some piece of iron, for example, to a large nail, cleaned to a shine. Lower the iron electrode into the jar so that it does not touch the bottom and walls. Figure out how to hang it yourself, it’s a simple thing. Pour a solution of alkali-caustic soda (handle with extreme caution!) or washing soda into the jar; The first option is better, but requires extreme care in work.

Since an alkali solution will be needed for experiments more than once, we will tell you here how to prepare it. Add washing soda Na 2 CO 3 to the slaked lime solution Ca(OH) 2 and boil the mixture. As a result of the reaction, caustic soda NaOH and calcium carbonate are formed, i.e. chalk, practically insoluble in water. This means that in the solution, which after cooling must be filtered, only alkali will remain. But let's return to the experience with the tin can. Soon, gas bubbles will begin to form on the iron electrode, and the tin from the tin will gradually go into solution. Well, what if you need to get not a solution containing tin, but the metal itself? Well, this is possible too. Remove the iron electrode from the solution and replace it with a carbon one. Here, an old, worn-out battery will help you again, with a carbon rod in its zinc cup. Remove it and connect the wire to the negative pole of your power source. Spongy tin will settle on the rod during electrolysis, and if the voltage is selected 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 place it inside the jar into 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 charcoal rod with sponge tin, put it in a porcelain cup or a clean metal can and hold it on the fire. Soon the tin will be fused into a dense ingot. Do not touch it or the jar until they have cooled!

Part of the sponge 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 will get a solution of tin chloride. Prepare such a solution with a concentration of approximately 7% and add, stirring, an alkali solution of a slightly higher concentration, about 10%. At first a white precipitate will form, but it will soon dissolve in the excess alkali. You have received a solution of sodium stannite - the same one that formed in the beginning when you began to dissolve tin from a jar. But if so, then the first part of the experiment - transferring the metal from the jar into the solution - can no longer be repeated, but proceed immediately to the 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 metal shoe polish can and heat over a flame. Once the lead has melted, carefully remove the jar from the heat by grasping the side of the jar with large, secure tweezers or pliers. Pour the melted lead into a plaster or metal mold, or simply into a sand hole - this way you will get a homemade lead casting. If you continue to calcinate molten lead 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 in the past to make paints.

Lead, like many other metals, interacts 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. Lead in it, although slowly, dissolves!

This paradox is explained by the fact that when interacting with hydrochloric acid, poorly soluble lead chloride PbCl 2 is formed. By 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 carry out 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 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 (careful!). 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, NaAlO 2 aluminate.

Aluminum oxide and hydroxide exhibit both basic and acidic properties, i.e. they react with both acids and alkalis. They are called amphoteric. Tin compounds, by the way, are also amphoteric; check it out for yourself, assuming you've already removed the tin from the tin.

There is a rule: the more active the metal, the more likely it is to oxidize and corrode. 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 it is kept in air and water, practically does not change. What is this - an exception to the rule?

Let's do an experiment. Fix a piece of aluminum wire in an inclined position over the flame of a gas burner or alcohol lamp so that the lower part of the wire is heated. At 660 °C this metal melts; it would seem that you would 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 containing molten metal. This case is made of aluminum oxide Al 2 O 3, a durable and very heat-resistant substance.

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

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

Let's try to analyze. Break off a piece of the coating from the old part and leave it in the air for several days so that it has time to become covered with a film of oxide, and then place it in a test tube with concentrated hydrochloric acid (handle with care! The acid should not get on your hands or clothes!). If it was nickel, then it will immediately begin to dissolve in the acid, forming the 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 the acid with the formation of chromium chloride CrCl 3. By removing this piece of coating from the acid with tweezers, rinsing it with water and air drying it, after two or three days you can observe the same effect again.

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, albeit 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 the air before the first experiment? After all, the chromium was already covered with a layer of oxide! And then, only the outer side was covered, and the inner side, facing the product, did not come into contact with oxygen in the air.

EXPERIMENTS WITH COPPER WIRE

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

Make a small spiral from a piece of copper wire and secure it in a wooden holder (you can leave a free end of sufficient length and wrap it around a regular pencil). Heat the coil in a flame. Its surface will be covered with a black coating of copper oxide CuO. If a blackened wire is dipped into 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 the salt CuCl 2.

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

Try finding an old copper object, say a candlestick. Scrape off some of the green residue and place it in a test tube. Close the neck of the test tube with a stopper 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 tube, causing the limewater to become cloudy. So it's carbon dioxide. What remains in the test tube is a black powder, which when dissolved in acid gives a blue solution. This powder, as you probably guess, is copper oxide.

So, we found out what components green plaque decomposes into. Its formula is written as follows: CuCO 3 * Cu (OH) 2 (basic copper carbonate). It forms on copper objects because there is always both carbon dioxide and water vapor in the air. The green coating is called patina. The same salt is found in nature - it is none other than the famous mineral malachite.

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

Pour ammonia into a test tube, heat the copper wire red-hot and lower it into the vial. The spiral will hiss and again become red and shiny. In an instant, a reaction will occur resulting in the formation of copper, water and nitrogen. If the experiment is repeated several times, the ammonia in the test tube will turn blue. Simultaneously with this reaction, another, so-called complexation reaction occurs - the same copper complex compound is formed, which previously allowed us to accurately identify 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 (even since those times when the science of chemistry was not even in sight). Copper and brass objects were cleaned with an ammonia solution, i.e., ammonia, to a shine. This, by the way, is what experienced housewives do now; for greater effect, ammonia is mixed with chalk, which mechanically scrubs away dirt and adsorbs contaminants from the solution.

Next experience. Pour a little ammonia into the 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). Using a hot copper spiral, touch the layer of substance covering the bottom of the test tube. The hissing will be heard again, and white smoke will rise up - these are the particles of ammonia evaporating, and the spiral will again sparkle with its pristine copper shine. A reaction occurred, as a result of which the same products were formed as in the previous experiment, and in addition copper chloride CuCl 2.

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

And the last experiment with a copper spiral. Pour a little cologne into the test tube (even better - pure alcohol) and again introduce the hot copper wire. In all likelihood, you can already imagine the result of the experiment: the wire was again cleared of the oxide film. This time a complex organic reaction occurred: 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 needs to be obtained from alcohol.

That's it for our first, introductory experiments. Now that you have, as they say, gotten your hands on in the experiment, and if you are doing experiments at home, you have probably created a certain supply of glassware and available reagents, it’s time to do more serious experiments. Let's take a look in the kitchen cabinet...

Hydrogen peroxide, which is the basis of our experience, is a very unstable compound. The 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 occurs very slowly. 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 3% hydrogen peroxide solution into a plastic bottle. This solution is sold in pharmacies as an antiseptic. Instead of peroxide, you can use bleach - they are also prepared on the basis of H2O2.

Hydrogen peroxide (as peroxide is also called) is dangerous for living beings. An enzyme called catalase is used to break down H2O2 into oxygen and water. Catalase is found in almost all living organisms, including the yeast we use in our experiments.

2. Add food coloring. It is better to use food paints - not because we are going to eat foam (this is not healthy in any case), 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/cm 3 . 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 best. The volume is approximately half the volume of peroxide, that is, 100 ml.

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

4. Dissolve the yeast in warm water using a separate glass. This is not so easy to do - the yeast will stick together in clumps. You need to patiently stir a tablespoon of yeast into 50 ml of water and then let it sit for five minutes. Determinedly pour the yeast solution into the hydrogen peroxide bottle and prepare to observe. If you're lucky, the reaction will be 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 foaming dishwashing detergent to the hydrogen peroxide solution.