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ENTERTAINING EXPERIENCES AT HOME
Attention, speed! Chemical experiments
Entertaining experiences at home / Chemistry experiments for children In chemical science there is a special area that studies the rates and mechanisms of various reactions - chemical kinetics. Although chemical theory can explain a lot, it is not yet possible to theoretically predict the rate of any reaction. It is studied experimentally, in the laboratory, and then they develop ways to change this speed. There are many reactions that are important for the industry, which are too slow, you need to be able to speed them up. Other reactions, on the contrary, have to be inhibited because they are harmful. In short, chemical kinetics is an experimental science. The validity of its laws can be verified by a few simple experiments. To begin with, let's make sure that the rate of the same reaction can really change, and quite significantly. (However, this can be assumed on the basis of not chemical, but life experience; for example, food in the cold deteriorates more slowly than in the heat, because at different temperatures the same biochemical reactions proceed at different speeds.) To check, repeat the experiment from the chapter "chemical clock", but this time change not the concentration of substances (this is already familiar to you), but the temperature. If both initial solutions - sodium sulfate and potassium iodate with sulfuric acid - are poured into ice water, then the time before the appearance of a blue color will take noticeably longer, than when using warm water.Note only that in very hot water no color appears at all, since the colored combination of iodine with starch is unstable. So, you found out by experience: the higher the concentration and temperature, the faster the reaction. But some reactions at first glance seem to be an exception to the rule. Here is an example. Pour acetic acid into a test tube to a height of 1-2 cm and drop a few pieces of zinc into it. Zinc must first be cleaned by immersing it for twenty seconds in a solution of hydrochloric acid and rinsing with water. Acetic acid is weak, and zinc dissolves in it very slowly - hydrogen bubbles are barely released. How to speed up the reaction? Heat up the solution. Right. Is it not possible otherwise? Let's do this: we will gradually add clean water to the test tube, mixing well each time. Keep a close eye on the bubbles. An amazing thing: the acid is already diluted by half, three times, and the reaction, instead of slowing down, is going faster and faster! If you put this experience in the circle, then replace the zinc with a small piece of magnesium shavings and do not process it with anything. With dilute acetic acid, magnesium reacts even more vigorously than zinc. Such an "exception" to the rule becomes clear if you study it well. Our experience with acetic acid is explained as follows. The rate at which zinc or magnesium reacts with an acid depends on the concentration of hydrogen ions in the solution. These ions are formed when any acid is dissolved in water. But when water is scarce, weak acetic acid is found in solution almost exclusively in the form of undissociated molecules. As it is diluted with water, more and more acetic acid molecules break down into ions, and the reaction proceeds faster. But if you add too much water, then the reaction will slow down again, for a different reason: due to strong dilution, the concentration of hydrogen ions will decrease again. 15% acetic acid reacts most rapidly with zinc. Of course, we have analyzed this experiment by no means in order to simply show how unusual chemical transformations can be. We wanted to draw your attention to this: to control the speed of a reaction, you must know how it goes. Any reaction begins with the fact that the molecules of substances collide with each other. Let's see how the reaction starts.
Take a not very wide glass tube several tens of centimeters long and pick up two corks to it. From the inside facing the tube, insert a small glass rod into both corks and wrap a piece of cotton wool around them. Moisten one piece with a few drops of concentrated hydrochloric acid, the other with concentrated ammonia solution. At the same time, insert the plugs with cotton wool into the tube from both ends. After a few minutes - depending on the length of the tube - in it, closer to the cotton wool with hydrochloric acid, a white ring of ammonium chloride NH will appear4Cl. Usually, during chemical reactions, the mixture is stirred to make the process go faster. We deliberately did not do this and did not even try to help the molecules meet - they moved on their own. Such independent movement of molecules in one or another medium is called diffusion. Evaporating from cotton wool, the molecules of both substances experienced billions of collisions per second with air molecules and with each other. And although the speed of molecules is very high, it is calculated in hundreds of meters per second, at 0 ° C and normal pressure, the free path, that is, the distance that a molecule manages to travel from one collision to another, is only about 0,0001 mm for these substances . That is why ammonia and hydrogen chloride (from hydrochloric acid) moved so slowly in the tube. Just as slowly, an odorous substance spreads through a room with still air. But why didn't the white ring appear in the middle of the tube? Because the ammonia molecules are smaller, they move through the air faster. If air is pumped out of the tube, then the molecules of ammonia and hydrogen chloride will meet in a fraction of a second - the mean free path of the molecules will increase significantly. We advise you to do a little research on your own to find out how gravity and temperature affect diffusion. To do this, position the tube vertically and obliquely, and also heat individual parts of it (including the place where ammonium chloride settles). Try to draw your own conclusions. Let's move on from gases to liquids. In them, diffusion proceeds even more slowly. Let's check it out experimentally. On a smooth and clean glass plate, drop a few drops of three liquids next to each other: in the middle - water, on the sides of it - solutions of soda and hydrochloric acid. Liquids before the start of the experiment should not come into contact. Then, very carefully, avoiding stirring, combine the solutions with a stick. Carbon dioxide should be released, but this will not happen immediately. And when the gas begins to be released, then its bubbles will be located along the boundary separating the areas of diffusion of acid and soda. Instead of soda and acid, you can take any two water-soluble substances that, when mixed, color or precipitate. However, in such experiments it is difficult to avoid liquid flows that distort the picture, so it is better to conduct experiments in thickened solutions. And you can thicken them with gelatin. Prepare a 4% gelatin solution by immersing it in hot water (do not boil!). Pour the hot solution into a test tube and, when it cools down, quickly, in one motion, insert a crystal of potassium permanganate, copper sulfate or another brightly colored and water-soluble substance into the center of the test tube with tweezers. Remove the tweezers immediately with a careful but quick movement. Within a few hours, a very beautiful diffusion pattern can be observed. The solute propagates in all directions at the same speed, forming a colored sphere. With a thickened solution, you can put another experiment. Pour the hot gelatin solution into two test tubes and add a little alkali solution to one and phenolphthalein to the other. When the contents of the test tubes harden, with tweezers quickly insert a piece of a phenolphthalein tablet into the center of the first test tube, and a lump of soda ash into the center of the second. In both cases, a crimson color will appear. But note: in the second test tube, the color spreads much faster. Hydroxide ions formed during the dissociation of alkali are much smaller and lighter than the complex organic phenolphthalein molecule, and therefore they move faster in solution. Now let's move on to solids. In reactions between them (or between a solid and a liquid or gas), molecules can collide only on the surface. The larger the interface, the faster the reaction. Let's make sure of this. Iron does not burn in air. However, this is true only for iron objects. For example, nails have a small surface of contact with air, the oxidation reaction is too slow. Iron filings react with oxygen much faster: in the cold they turn into rust earlier, and in a flame they can catch fire. The smallest grains can flare up without heating at all. Such iron is called pyrophoric. It cannot be planed even with the smallest file, so it is obtained chemically, for example, by decomposing the salt of oxalic acid - iron oxalate. Mix aqueous solutions of some ferrous salt, such as ferrous sulfate, and oxalic acid or its soluble salt. Filter the yellow precipitate of iron oxalate and fill the test tube with it no more than a fifth of the volume. Heat the substance in the flame of the burner, while holding the test tube horizontally or slightly inclined, with the hole down and away from you. Remove any droplets of water that come out with a filter paper or cotton wool. When the oxalate has decomposed and turned into a black powder, close the vial and refrigerate it. Gradually and very carefully pour the contents of the test tube onto a metal or asbestos sheet: the powder will flare up with bright sparks. The experience is especially effective in a darkened room. Important warning: pyrophoric iron must not be stored, it can cause a fire! At the end of the experiment, be sure to ignite the powder in air or treat with acid so that there are no unburned particles left - they can ignite spontaneously. Next, we investigate how the size of the surface of a solid affects the rate of its reaction with a liquid. Take two identical pieces of chalk and grind one of them into powder. Place both samples in tubes and fill with equal volumes of hydrochloric acid. Finely divided chalk, as one would expect, will dissolve much faster. Place another piece of chalk in a test tube with sulfuric acid. The energetic reaction that began at first subsides, and then stops altogether. From what? After all, sulfuric acid is not weaker than hydrochloric acid ... When chalk reacts with hydrochloric acid, calcium chloride CaCl is formed.2 which is easily soluble in water and does not interfere with the flow of new portions of acid to the surface of the chalk. When reacted with sulfuric acid, calcium sulfate CaSO is obtained.4, and it is very poorly soluble in water, remains on the surface of the chalk and closes it. In order for the reaction to proceed further, it is necessary to clean the surface of the chalk from time to time or turn it into powder in advance. Knowledge of such process details is very important for chemical engineering. And one more experience. Mix in a mortar and mortar two solid substances that give colored reaction products: lead nitrate and potassium iodide, ferrous sulfate and red blood salt, etc. - and rub the mixture with a pestle. Gradually, as the mixture is rubbed, the mixture will begin to color, as the interaction surface of the substances increases. If you pour a little water on the mixture, then an intense color will immediately appear - after all, the molecules move much easier in the solution. And in conclusion of experiments on kinetics, we will set up a quantitative experiment; the only tool you'll need is a stopwatch or watch with a second hand. Prepare 0,5 l of a 3% solution of sulfuric acid (pouring acid into water!) And the same amount of a 12% solution of sodium thiosulfate. Before dissolving the thiosulfate, add a few drops of ammonia to the water. On two cylindrical bottles (glasses, stacks) with a capacity of 100 ml, mark at level 50; 25; 12,5 and 37,5 ml, sequentially dividing the height in half. Mark the bottles and pour the prepared solutions into them up to the upper marks (50 ml). Place an ordinary thin glass with a capacity of 200 or 250 ml on dark paper and pour a solution of thiosulfate into it, and then acid. Immediately note the time and stir the mixture for one to two seconds. In order not to break the glass, it is better to use a wooden stick. As soon as the solution begins to turn cloudy, record the time elapsed since the start of the reaction. It is convenient to carry out the experiment together: one monitors the clock, the other drains the solutions and signals cloudiness. Wash the glass and repeat the experiment three more times; pour the thiosulfate solution into a glass up to the third (37,5), second (25) and first (12,5 ml) marks, adding water each time to the upper division. The amount of acid in all experiments remains constant, and the total volume of the reacting mixture is always 100 ml. Now draw a graph: how the reaction rate depends on the concentration of thiosulfate. It is convenient to express the concentration in arbitrary units: 1, 2, 3 and 4. Put them on the x-axis. But how do you calculate the rate of a reaction? This cannot be done exactly, if only because we determine the moment of turbidity by eye, to a certain extent subjectively. In addition, the turbidity only shows that the smallest particles of sulfur that are released during the reaction have reached such a size that they can be seen. And yet, for lack of a better way, let's take the beginning of turbidity as the end of the reaction (which, by the way, is not very far from the truth). Let's make one more assumption: the reaction rate is inversely proportional to its duration. If the reaction took 10 seconds, then we will assume that the rate is 0,1. Plot the velocities on the y-axis. Four experiments gave four points, the fifth - the origin. All five points will be located approximately on one straight line. Her equation is written like this: v == k [Na2S2O3] where v- is the rate of the reaction, square brackets are the designation of concentration accepted in chemical kinetics, and k- the rate constant, which is easy to find from the graph. But the reaction rate must also depend on the concentration of sulfuric acid. By leaving the amount of thiosulfate unchanged and diluting the sulfuric acid, check how the reaction rate changes. Oddly enough, it does not change! Such cases are not uncommon. In our experiment, a complex reaction takes place, and its product, sulfur, is not immediately released during direct collisions of thiosulfate and acid molecules. In general, there are not so many reactions where products are obtained immediately. In complex sequential reactions Some stages are slower than others. In our case, the latter, in which sulfur is formed. It was her speed that we, in fact, measured. Author: Olgin O.M.
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