Wet relationship - part 1
Inorganic compounds are usually not associated with moisture, while organic compounds are vice versa. After all, the former are dry rocks, and the latter come from aquatic living organisms. However, widespread associations have little to do with reality. In this case, it is similar: water can be squeezed out of stones, and organic compounds can be very dry.
Water is a ubiquitous substance on Earth, and it is not surprising that it can be found in other chemical compounds as well. Sometimes it is loosely connected with them, enclosed within them, manifests itself in a latent form or openly builds the structure of crystals.
First things first. At the beginning…
… Moisture
Many chemical compounds tend to absorb water from their environment - for example, the well-known table salt, which often clumps together in the steamy and humid atmosphere of the kitchen. Such substances are hygroscopic and the moisture they cause hygroscopic water. However, table salt requires a high enough relative humidity (see box: How much water is in the air?) to bind the water vapor. Meanwhile, in the desert there are substances that can absorb water from the environment.
How much water is in the air?
Absolute humidity is the amount of water vapor contained in a unit volume of air at a given temperature. For example, at 0°С in 1 m3 In the air there can be a maximum (so that there is no condensation) of about 5 g of water, at 20 ° C - about 17 g of water, and at 40 ° C - more than 50 g. In a warm kitchen or bathroom, this is therefore quite wet.
Relative humidity is the ratio of the amount of water vapor per unit volume of air to the maximum amount at a given temperature (expressed as a percentage).
The next experiment will require sodium NaOH or potassium hydroxide KOH. Place a compound tablet (as they are sold) on a watch glass and leave in the air for a while. Soon you will notice that the lozenge begins to be covered with drops of liquid, and then spread. This is the effect of the hygroscopicity of NaOH or KOH. By placing the samples in different rooms of the house, you can compare the relative humidity of these places (1).
1. Precipitation of NaOH on a watch glass (left) and the same precipitate after a few hours in air (right).
2. Laboratory desiccator with silicone gel (photo: Wikimedia/Hgrobe)
Chemists, and not only them, solve the problem of the moisture content of a substance. Hygroscopic water it is an unpleasant contamination by a chemical compound, and its content, moreover, is unstable. This fact makes it difficult to weigh the amount of reagent required for the reaction. The solution, of course, is to dry the substance. On an industrial scale, this happens in heated chambers, that is, in an enlarged version of a home oven.
In laboratories, in addition to electric dryers (again, ovens), exykatory (also for storage of already dried reagents). These are glass vessels, tightly closed, at the bottom of which there is a highly hygroscopic substance (2). Its job is to absorb moisture from the dried compound and keep the humidity inside the desiccator low.
Examples of desiccants: Anhydrous CaCl salts.2 i MgSO4, phosphorus (V) oxides P4O10 and calcium CaO and silica gel (silica gel). You will also find the latter in the form of desiccant sachets placed in industrial and food packaging (3).
3. Silicone gel to protect food and industrial products from moisture.
Many dehumidifiers can be regenerated if they absorb too much water - just warm them up.
There is also chemical contamination. bottled water. It penetrates into the crystals during their rapid growth and creates spaces filled with the solution from which the crystal formed, surrounded by a solid. You can get rid of the liquid bubbles in the crystal by dissolving the compound and recrystallizing it, but this time under conditions that slow the growth of the crystal. Then the molecules will “neatly” settle down in the crystal lattice, leaving no gaps.
hidden water
In some compounds, water exists in a latent form, but the chemist is able to extract it from them. It can be assumed that you will release water from any oxygen-hydrogen compound under the right conditions. You will make it give up water by heating or by the action of another substance which strongly absorbs water. Water in such a relationship constitutional water. Try both chemical dehydration methods.
4. Water vapor condenses in the test tube when chemicals are dehydrated.
Pour a little baking soda into the test tube, i.e. sodium bicarbonate NaHCO.3. You can get it at the grocery store, and it's used in the kitchen, for example. as a leavening agent for baking (but also has many other uses).
Place the test tube in the flame of the burner at an angle of approximately 45° with the exit opening facing you. This is one of the principles of laboratory hygiene and safety - this is how you protect yourself in the event of a sudden release of a heated substance from a test tube.
Heating is not necessarily strong, the reaction will begin at 60 ° C (a methylated spirit burner or even a candle is enough). Keep an eye on the top of the vessel. If the tube is long enough, drops of liquid will begin to collect at the outlet (4). If you do not see them, place a cold watch glass over the test tube outlet - water vapor released during the decomposition of baking soda condenses on it (the symbol D above the arrow indicates the heating of the substance):
5. Black hose comes out of the glass.
The second gaseous product, carbon dioxide, can be detected using lime water, i.e. saturated solution calcium hydroxide Sa (ON)2. Its turbidity caused by precipitation of calcium carbonate is indicative of the presence of CO2. It is enough to take a drop of the solution on a baguette and place it on the end of the test tube. If you don't have calcium hydroxide, make lime water by adding a NaOH solution to any water-soluble calcium salt solution.
In the next experiment, you will use the next kitchen reagent - ordinary sugar, that is, sucrose C.12H22O11. You will also need a concentrated solution of sulfuric acid H2SO4.
I immediately remind you of the rules for working with this dangerous reagent: rubber gloves and goggles are required, and the experiment is carried out on a plastic tray or plastic wrap.
Pour sugar into a small beaker half as much as the vessel is filled. Now pour in a solution of sulfuric acid in an amount equal to half the poured sugar. Stir the contents with a glass rod so that the acid is evenly distributed throughout the volume. Nothing happens for a while, but suddenly the sugar begins to darken, then turns black, and finally begins to "leave" the vessel.
A porous black mass, no longer looking like white sugar, crawls out of the glass like a snake from a fakirs' basket. The whole thing warms up, clouds of water vapor are visible and even a hiss is heard (this is also water vapor escaping from the cracks).
The experience is attractive, from the category of the so-called. chemical hoses (5). The hygroscopicity of a concentrated solution of H is responsible for the observed effects.2SO4. It is so large that water enters the solution from other substances, in this case sucrose:
Residues of sugar dehydration are saturated with water vapor (remember that when mixing concentrated H2SO4 a lot of heat is released with water), which causes a significant increase in their volume and the effect of lifting the mass from the glass.
Trapped in a crystal
6. Heating of crystalline copper sulfate (II) in a test tube. Partial dehydration of the compound is visible.
And another kind of water contained in chemicals. This time it appears explicitly (unlike constitutional water), and its amount is strictly defined (and not arbitrary, as in the case of hygroscopic water). This water of crystallizationwhat gives color to the crystals - when removed, they disintegrate into an amorphous powder (which you will see experimentally, as befits a chemist).
Stock up on blue crystals of hydrated copper(II) sulfate CuSO4× 5h2Oh, one of the most popular laboratory reagents. Pour a small amount of small crystals into a test tube or evaporator (the second method is better, but in the case of a small amount of the compound, a test tube can also be used; more on that in a month). Gently start heating over the burner flame (a denatured alcohol lamp will suffice).
Frequently shake the tube away from you, or stir the baguette in the evaporator placed in the tripod handle (do not lean over the dishes). As the temperature rises, the color of the salt begins to fade, until finally it becomes almost white. In this case, drops of liquid collect in the upper part of the test tube. This is the water removed from the salt crystals (heating them in an evaporator will reveal the water by placing a cold watch glass over the vessel), which has meanwhile disintegrated into a powder (6). The dehydration of the compound occurs in stages:
A further increase in temperature above 650°C causes decomposition of the anhydrous salt. White powder anhydrous CuSO4 store in a tightly screwed container (you can put a moisture-absorbing bag in it).
You may ask: how do we know that dehydration occurs as described by the equations? Or why relationships follow this pattern? You will work on determining the amount of water in this salt next month, now I will answer the first question. The method by which we can observe the change in the mass of a substance with increasing temperature is called thermogravimetric analysis. The test substance is placed on a pallet, the so-called thermal balance, and heated, reading the weight changes.
Of course, today thermobalances record the data themselves, at the same time drawing the corresponding graph (7). The shape of the curve of the graph shows at what temperature "something" happens, for example, a volatile substance is released from the compound (loss of weight) or it combines with a gas in the air (then the mass increases). The change in mass allows you to determine what and in what quantity has decreased or increased.
7. Graph of the thermogravimetric curve of crystalline copper(II) sulfate.
Hydrated CuSO4 it has almost the same color as its aqueous solution. This is not a coincidence. Cu ion in solution2+ is surrounded by six water molecules, and in the crystal - by four, lying at the corners of the square, the center of which it is. Above and below the metal ion are sulfate anions, each of which "serves" two adjacent cations (so the stoichiometry is correct). But where is the fifth water molecule? It lies between one of the sulfate ions and a water molecule in a belt surrounding the copper(II) ion.
And again, the inquisitive reader will ask: how do you know this? This time from images of crystals obtained by irradiating them with X-rays. However, explaining why an anhydrous compound is white and a hydrated compound is blue is advanced chemistry. It's time for her to study.
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