Soda, volcano, Venus, refrigerator - what do they have in common? Carbon dioxide. We have collected for you the most interesting information about one of the most important chemical compounds on earth.
What is carbon dioxide
Carbon dioxide is known mainly in its gaseous state, i. as carbon dioxide with simple chemical formula CO2. In this form, it exists under normal conditions - at atmospheric pressure and "normal" temperatures. But at increased pressure, over 5,850 kPa (such, for example, the pressure at a sea depth of about 600 m), this gas turns into a liquid. And with strong cooling (minus 78.5 ° C), it crystallizes and becomes the so-called dry ice, which is widely used in trade for storing frozen foods in refrigerators.
Liquid carbon dioxide and dry ice are produced and used in human activities, but these forms are unstable and break down easily.
But gaseous carbon dioxide is ubiquitous: it is released during the respiration of animals and plants and is an important part of chemical composition atmosphere and ocean.
Properties of carbon dioxide
Carbon dioxide CO2 is colorless and odorless. Under normal conditions, it has no taste. However, when inhaling high concentrations of carbon dioxide, a sour taste can be felt in the mouth, caused by the fact that carbon dioxide dissolves on mucous membranes and in saliva, forming a weak solution of carbonic acid.
By the way, it is the ability of carbon dioxide to dissolve in water that is used to make sparkling waters. Bubbles of lemonade - the same carbon dioxide. The first apparatus for saturating water with CO2 was invented as early as 1770, and already in 1783, the enterprising Swiss Jacob Schwepp began the industrial production of soda (the Schweppes trademark still exists).
Carbon dioxide is 1.5 times heavier than air, so it tends to “settle” in its lower layers if the room is poorly ventilated. The “dog cave” effect is known, where CO2 is released directly from the ground and accumulates at a height of about half a meter. An adult, getting into such a cave, at the height of his height does not feel an excess of carbon dioxide, but dogs find themselves right in a thick layer of carbon dioxide and are poisoned.
CO2 does not support combustion, so it is used in fire extinguishers and fire suppression systems. The trick with extinguishing a burning candle with the contents of an allegedly empty glass (but in fact with carbon dioxide) is based precisely on this property of carbon dioxide.
Carbon dioxide in nature: natural sources
Carbon dioxide is produced in nature from various sources:
- Breathing of animals and plants.
Every schoolchild knows that plants absorb carbon dioxide CO2 from the air and use it in photosynthesis. Some housewives are trying to atone for shortcomings with an abundance of indoor plants. However, plants not only absorb but also release carbon dioxide in the absence of light as part of the respiration process. Therefore, a jungle in a poorly ventilated bedroom is not a good idea: at night, CO2 levels will rise even more. - Volcanic activity.
Carbon dioxide is part of volcanic gases. In areas with high volcanic activity CO2 can be emitted directly from the ground - from cracks and fissures called mofets. The concentration of carbon dioxide in mofet valleys is so high that many small animals die when they get there. - Decomposition organic matter.
Carbon dioxide is formed during combustion and decay of organic matter. Volumetric natural emissions of carbon dioxide accompany forest fires.
Carbon dioxide is "stored" in nature in the form of carbon compounds in minerals: coal, oil, peat, limestone. Huge reserves of CO2 are found in dissolved form in the world's oceans.
The release of carbon dioxide from an open reservoir can lead to a limnological catastrophe, as happened, for example, in 1984 and 1986. in lakes Manun and Nyos in Cameroon. Both lakes were formed on the site of volcanic craters - now they are extinct, but in the depths, volcanic magma still emits carbon dioxide, which rises to the waters of the lakes and dissolves in them. As a result of a number of climatic and geological processes, the concentration of carbon dioxide in the waters exceeded the critical value. Was released into the atmosphere great amount carbon dioxide, which, like an avalanche, descended down the mountain slopes. About 1,800 people became victims of limnological disasters on the Cameroonian lakes.
Artificial sources of carbon dioxide
The main anthropogenic sources of carbon dioxide are:
- industrial emissions associated with combustion processes;
- automobile transport.
Despite the fact that the share of environmentally friendly transport in the world is growing, the vast majority of the world's population will not soon be able (or willing) to switch to new cars.
Active deforestation for industrial purposes also leads to an increase in the concentration of carbon dioxide CO2 in the air.
CO2 is one of the end products of metabolism (the breakdown of glucose and fats). It is secreted in the tissues and carried by hemoglobin to the lungs, through which it is exhaled. In the air exhaled by a person, there is about 4.5% carbon dioxide (45,000 ppm) - 60-110 times more than in the inhaled air.
Carbon dioxide plays an important role in the regulation of blood supply and respiration. An increase in the level of CO2 in the blood causes the capillaries to expand, allowing more blood to pass through, which delivers oxygen to the tissues and removes carbon dioxide.
The respiratory system is also stimulated by an increase in carbon dioxide, and not by a lack of oxygen, as it might seem. In fact, the lack of oxygen is not felt by the body for a long time, and it is quite possible that in rarefied air a person will lose consciousness before he feels a lack of air. The stimulating property of CO2 is used in artificial respiration devices: there, carbon dioxide is mixed with oxygen to "start" the respiratory system.
Carbon dioxide and us: why is CO2 dangerous?
Carbon dioxide is as essential to the human body as oxygen. But just like with oxygen, an excess of carbon dioxide harms our well-being.
A high concentration of CO2 in the air leads to intoxication of the body and causes a state of hypercapnia. In hypercapnia, a person experiences difficulty breathing, nausea, headache, and may even pass out. If the carbon dioxide content does not decrease, then the turn comes - oxygen starvation. The fact is that both carbon dioxide and oxygen move around the body on the same "transport" - hemoglobin. Normally, they "travel" together, attaching to different places on the hemoglobin molecule. However, an increased concentration of carbon dioxide in the blood reduces the ability of oxygen to bind to hemoglobin. The amount of oxygen in the blood decreases and hypoxia occurs.
Such unhealthy consequences for the body occur when inhaling air with a CO2 content of more than 5,000 ppm (this can be the air in mines, for example). To be fair, in ordinary life we practically do not encounter such air. However, even a much lower concentration of carbon dioxide is not good for health.
According to the findings of some, already 1,000 ppm CO2 causes fatigue and headache in half of the subjects. Many people begin to feel closeness and discomfort even earlier. With a further increase in the concentration of carbon dioxide to 1,500 - 2,500 ppm, the brain is "lazy" to take the initiative, process information and make decisions.
And if the level of 5,000 ppm is almost impossible in Everyday life, then 1,000 and even 2,500 ppm can easily be part of reality modern man. Ours showed that in sparsely ventilated classrooms, CO2 levels stay above 1,500 ppm most of the time, and sometimes jump above 2,000 ppm. There is every reason to believe that the situation is similar in many offices and even apartments.
Physiologists consider 800 ppm as a safe level of carbon dioxide for human well-being.
Another study found a link between CO2 levels and oxidative stress: the higher the level of carbon dioxide, the more we suffer from, which destroys the cells of our body.
Carbon dioxide in the earth's atmosphere
In the atmosphere of our planet, there is only about 0.04% CO2 (this is approximately 400 ppm), and more recently it was even less: carbon dioxide crossed the mark of 400 ppm only in the fall of 2016. Scientists attribute the rise in the level of CO2 in the atmosphere to industrialization: in the middle of the 18th century, on the eve of the industrial revolution, it was only about 270 ppm.
Let's imagine the following situation:
You work in a lab and decide to do an experiment. To do this, you opened the cabinet with reagents and suddenly saw the following picture on one of the shelves. Two jars of reagents had their labels peeled off, which were safely left lying nearby. At the same time, it is no longer possible to determine exactly which jar corresponds to which label, and the external signs of the substances by which they could be distinguished are the same.
In this case, the problem can be solved using the so-called qualitative reactions.
Qualitative reactions called such reactions that allow you to distinguish one substance from another, as well as to find out qualitative composition unknown substances.
For example, it is known that the cations of some metals, when their salts are added to the burner flame, color it in a certain color:
This method can only work if the substances to be distinguished change the color of the flame in different ways, or one of them does not change color at all.
But, let's say, as luck would have it, the substances you determine do not color the color of the flame, or color it in the same color.
In these cases, it will be necessary to distinguish substances using other reagents.
In what case can we distinguish one substance from another with the help of any reagent?
There are two options:
- One substance reacts with the added reagent, while the other does not. At the same time, it must be clearly visible that the reaction of one of the starting substances with the added reagent has really passed, that is, some external sign of it is observed - a precipitate has formed, a gas has been released, a color change has occurred, etc.
For example, it is impossible to distinguish water from sodium hydroxide solution using of hydrochloric acid, despite the fact that alkalis with acids react perfectly:
NaOH + HCl \u003d NaCl + H 2 O
This is due to the absence of any external signs reactions. A transparent colorless solution of hydrochloric acid, when mixed with a colorless hydroxide solution, forms the same transparent solution:
But on the other hand, water can be distinguished from an aqueous solution of alkali, for example, using a solution of magnesium chloride - a white precipitate forms in this reaction:
2NaOH + MgCl 2 = Mg(OH) 2 ↓+ 2NaCl
2) Substances can also be distinguished from each other if they both react with the added reagent, but do so in different ways.
For example, a solution of sodium carbonate can be distinguished from a solution of silver nitrate using a solution of hydrochloric acid.
hydrochloric acid reacts with sodium carbonate to release a colorless, odorless gas - carbon dioxide (CO 2):
2HCl + Na 2 CO 3 \u003d 2NaCl + H 2 O + CO 2
and with silver nitrate to form a white cheesy precipitate AgCl
HCl + AgNO 3 \u003d HNO 3 + AgCl ↓
The tables below show various options specific ion detection:
Qualitative reactions to cations
| Cation | Reagent | Sign of reaction |
| Ba 2+ | SO 4 2- |
Ba 2+ + SO 4 2- \u003d BaSO 4 ↓ |
| Cu2+ | 1) Precipitation of blue color: Cu 2+ + 2OH - \u003d Cu (OH) 2 ↓ 2) Precipitation of black color: Cu 2+ + S 2- \u003d CuS ↓ |
|
| Pb 2+ | S2- | Precipitation of black color: Pb 2+ + S 2- = PbS↓ |
| Ag+ | Cl- |
Precipitation of a white precipitate, insoluble in HNO 3, but soluble in ammonia NH 3 H 2 O: Ag + + Cl − → AgCl↓ |
| Fe2+ |
2) Potassium hexacyanoferrate (III) (red blood salt) K 3 |
1) Precipitation of a white precipitate that turns green in air: Fe 2+ + 2OH - \u003d Fe (OH) 2 ↓ 2) Precipitation of a blue precipitate (turnbull blue): K + + Fe 2+ + 3- = KFe↓ |
| Fe3+ |
2) Potassium hexacyanoferrate (II) (yellow blood salt) K 4 3) Rhodanide ion SCN − |
1) Precipitation of brown color: Fe 3+ + 3OH - \u003d Fe (OH) 3 ↓ 2) Precipitation of a blue precipitate (Prussian blue): K + + Fe 3+ + 4- = KFe↓ 3) The appearance of intense red (blood red) staining: Fe 3+ + 3SCN - = Fe(SCN) 3 |
| Al 3+ | Alkali (hydroxide amphoteric properties) |
Precipitation of a white precipitate of aluminum hydroxide when a small amount of alkali is added: OH - + Al 3+ \u003d Al (OH) 3 and its dissolution upon further addition: Al(OH) 3 + NaOH = Na |
| NH4+ | OH − , heating | Emission of gas with a pungent odor: NH 4 + + OH - \u003d NH 3 + H 2 O Blue wet litmus paper |
| H+ (acid environment) |
Indicators: − litmus − methyl orange |
Red staining |
Qualitative reactions to anions
| Anion | Impact or reagent | Reaction sign. Reaction equation |
| SO 4 2- | Ba 2+ |
Precipitation of a white precipitate, insoluble in acids: Ba 2+ + SO 4 2- \u003d BaSO 4 ↓ |
| NO 3 - |
1) Add H 2 SO 4 (conc.) and Cu, heat 2) A mixture of H 2 SO 4 + FeSO 4 |
1) Solution formation of blue color containing Cu 2+ ions, brown gas evolution (NO 2) 2) The appearance of the color of nitroso-iron sulfate (II) 2+. Violet to brown color (brown ring reaction) |
| PO 4 3- | Ag+ |
Precipitation of a light yellow precipitate in a neutral medium: 3Ag + + PO 4 3- = Ag 3 PO 4 ↓ |
| CrO 4 2- | Ba 2+ |
Precipitation of a yellow precipitate, insoluble in acetic acid, but soluble in HCl: Ba 2+ + CrO 4 2- = BaCrO 4 ↓ |
| S2- | Pb 2+ |
Black precipitation: Pb 2+ + S 2- = PbS↓ |
| CO 3 2- |
1) Precipitation of a white precipitate, soluble in acids: Ca 2+ + CO 3 2- \u003d CaCO 3 ↓ 2) Emission of a colorless gas ("boiling"), causing the lime water to become cloudy: CO 3 2- + 2H + = CO 2 + H 2 O |
|
| CO2 | Lime water Ca(OH) 2 |
Precipitation of a white precipitate and its dissolution upon further passage of CO 2: Ca(OH) 2 + CO 2 = CaCO 3 ↓ + H 2 O CaCO 3 + CO 2 + H 2 O \u003d Ca (HCO 3) 2 |
| SO 3 2- | H+ |
SO 2 gas evolution with a characteristic pungent odor (SO 2): 2H + + SO 3 2- \u003d H 2 O + SO 2 |
| F- | Ca2+ |
Precipitation of a white precipitate: Ca 2+ + 2F - = CaF 2 ↓ |
| Cl- | Ag+ |
Precipitation of a white cheesy precipitate, insoluble in HNO 3 but soluble in NH 3 H 2 O (conc.): Ag + + Cl - = AgCl↓ AgCl + 2(NH 3 H 2 O) =) Read also:
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