Oxygen (O) is a non-metallic chemical element of group 16 (VIa) of the periodic table. It is a colorless, odorless and tasteless gas that is essential for living organisms – animals that turn it into carbon dioxide and plants that use CO2 as a carbon source and return O 2 into the atmosphere. Oxygen forms compounds by reacting with almost any other element, and also displaces chemical elements from bonding with each other. In many cases, these processes are accompanied by the release of heat and light. The most important oxygen compound is water.
Discovery history
In 1772, Swedish chemist Carl Wilhelm Scheele first demonstrated oxygen by heating potassium nitrate, mercury oxide, and many other substances. Independently of him, in 1774, the English chemist Joseph Priestley discovered this chemical element by thermal decomposition of mercury oxide and published his findings in the same year, three years before publication. Scheele. In 1775-1780, the French chemist Antoine Lavoisier interpreted the role of oxygen in respiration and combustion, rejecting the phlogiston theory generally accepted at the time. He noted its tendency to form acids when combined with various substances and named the element oxygène, which in Greek means "producing acid".
Prevalence
What is oxygen? Making up 46% of the mass of the earth's crust, it is its most common element. The amount of oxygen in the atmosphere is 21% by volume, and by weight in sea water it is 89%.
In rocks, the element is combined with metals and non-metals in the form of oxides, which are acidic (for example, sulfur, carbon, aluminum and phosphorus) or basic (s alts of calcium, magnesium and iron), and as s alt-like compounds that can be be considered as formed from acidic and basic oxides such as sulfates, carbonates, silicates, aluminates and phosphates. Although they are numerous, these solids cannot serve as sources of oxygen, since breaking the bond of an element with metal atoms is too energy-consuming.
Features
If the temperature of oxygen is below -183 °C, then it becomes a pale blue liquid, and at -218 °C - solid. Pure O2 is 1.1 times heavier than air.
During respiration, animals and some bacteria consume oxygen from the atmosphere and return carbon dioxide, while during photosynthesis, green plants in the presence of sunlight absorb carbon dioxide and release free oxygen. Almostall O2 in the atmosphere is produced by photosynthesis.
At 20 °C, about 3 volume parts of oxygen dissolve in 100 parts of fresh water, slightly less in sea water. This is necessary for the breathing of fish and other marine life.
Natural oxygen is a mixture of three stable isotopes: 16O (99.759%), 17O (0.037 %) and 18O (0.204%). Several artificially produced radioactive isotopes are known. The longest-lived of these is 15O (with a half-life of 124 s), which is used to study respiration in mammals.
Allotropes
A clearer idea of what oxygen is, allows you to get its two allotropic forms, diatomic (O2) and triatomic (O3, ozone). The properties of the diatomic form suggest that six electrons bind the atoms and two remain unpaired, causing oxygen paramagnetism. The three atoms in the ozone molecule are not in a straight line.
Ozone can be produced according to the equation: 3O2 → 2O3.
The process is endothermic (requires energy); the conversion of ozone back into diatomic oxygen is facilitated by the presence of transition metals or their oxides. Pure oxygen is converted into ozone by a glowing electric discharge. The reaction also occurs upon absorption of ultraviolet light with a wavelength of about 250 nm. The occurrence of this process in the upper atmosphere eliminates radiation that could causedamage to life on the Earth's surface. The pungent smell of ozone is present in enclosed spaces with sparking electrical equipment such as generators. It is a light blue gas. Its density is 1.658 times that of air and it has a boiling point of -112°C at atmospheric pressure.
Ozone is a strong oxidizing agent, capable of converting sulfur dioxide to trioxide, sulfide to sulfate, iodide to iodine (providing an analytical method for evaluating it), and many organic compounds to oxygenated derivatives such as aldehydes and acids. The conversion of hydrocarbons from automobile exhaust gases into these acids and aldehydes by ozone is the cause of smog. In industry, ozone is used as a chemical agent, disinfectant, waste water treatment, water purification and fabric bleaching.
Getting Methods
The way oxygen is produced depends on how much gas is required. Laboratory methods are as follows:
1. Thermal decomposition of some s alts such as potassium chlorate or potassium nitrate:
- 2KClO3 → 2KCl + 3O2.
- 2KNO3 → 2KNO2 + O2.
The decomposition of potassium chlorate is catalyzed by transition metal oxides. Manganese dioxide (pyrolusite, MnO2) is often used for this. The catalyst lowers the temperature needed to evolve oxygen from 400 to 250°C.
2. Temperature decomposition of metal oxides:
- 2HgO → 2Hg +O2.
- 2Ag2O → 4Ag + O2.
Scheele and Priestley used a compound (oxide) of oxygen and mercury (II) to obtain this chemical element.
3. Thermal decomposition of metal peroxides or hydrogen peroxide:
- 2BaO + O2 → 2BaO2.
- 2BaO2 → 2BaO +O2.
- BaO2 + H2SO4 → H2 O2 + BaSO4.
- 2H2O2 → 2H2O +O 2.
The first industrial methods for separating oxygen from the atmosphere or for producing hydrogen peroxide depended on the formation of barium peroxide from the oxide.
4. Electrolysis of water with small impurities of s alts or acids, which provide the conductivity of electric current:
2H2O → 2H2 + O2
Industrial production
If it is necessary to obtain large volumes of oxygen, fractional distillation of liquid air is used. Of the major constituents of air, it has the highest boiling point and is therefore less volatile than nitrogen and argon. The process uses cooling of the gas as it expands. The main steps of the operation are as follows:
- air is filtered to remove particulate matter;
- moisture and carbon dioxide are removed by absorption into alkali;
- air is compressed and the heat of compression is removed by normal cooling procedures;
- then it enters the coil located incamera;
- part of the compressed gas (at a pressure of about 200 atm) expands in the chamber, cooling the coil;
- expanded gas is returned to the compressor and goes through several stages of subsequent expansion and compression, resulting in the air becoming liquid at -196 °C;
- liquid is heated to distill the first light inert gases, then nitrogen, and liquid oxygen remains. Multiple fractionation produces a product pure enough (99.5%) for most industrial purposes.
Industrial use
Metallurgy is the largest consumer of pure oxygen for the production of high-carbon steel: get rid of carbon dioxide and other non-metal impurities faster and easier than using air.
Oxygen wastewater treatment holds promise for treating liquid effluent more efficiently than other chemical processes. Waste incineration in closed systems using pure O2.
. is becoming increasingly important
The so-called rocket oxidizer is liquid oxygen. Pure O2 Used in submarines and diving bells.
In the chemical industry, oxygen has replaced normal air in the production of substances such as acetylene, ethylene oxide and methanol. Medical applications include the use of the gas in oxygen chambers, inhalers, and baby incubators. An oxygen-enriched anesthetic gas provides life support during general anesthesia. Without this chemical element, a number ofindustries using melting furnaces. That's what oxygen is.
Chemical properties and reactions
The high electronegativity and electron affinity of oxygen are typical of elements that exhibit non-metallic properties. All oxygen compounds have a negative oxidation state. When two orbitals are filled with electrons, an O2- ion is formed. In peroxides (O22-) each atom is assumed to have a charge of -1. This property of accepting electrons by total or partial transfer defines an oxidizing agent. When such an agent reacts with an electron donor substance, its own oxidation state is lowered. The change (decrease) in the oxidation state of oxygen from zero to -2 is called reduction.
Under normal conditions, the element forms diatomic and triatomic compounds. In addition, there are highly unstable four-atom molecules. In the diatomic form, two unpaired electrons are located in nonbonding orbitals. This is confirmed by the paramagnetic behavior of the gas.
The intense reactivity of ozone is sometimes explained by the assumption that one of the three atoms is in an "atomic" state. Entering the reaction, this atom dissociates from O3, leaving molecular oxygen.
The O2 molecule is weakly reactive at normal ambient temperatures and pressures. Atomic oxygen is much more active. The dissociation energy (O2 → 2O) is significant andis 117.2 kcal per mole.
Connections
With non-metals such as hydrogen, carbon and sulfur, oxygen forms a wide range of covalently bonded compounds, including oxides of non-metals such as water (H2O), sulfur dioxide (SO2) and carbon dioxide (CO2); organic compounds such as alcohols, aldehydes and carboxylic acids; common acids such as carbonic (H2CO3), sulfuric (H2SO 4) and nitrogen (HNO3); and corresponding s alts such as sodium sulfate (Na2SO4), sodium carbonate (Na2 CO3) and sodium nitrate (NaNO3). Oxygen is present in the form of the O2- ion in the crystal structure of solid metal oxides, such as the compound (oxide) of oxygen and calcium CaO. Metal superoxides (KO2) contain the O2- ion, while metal peroxides (BaO 2), contain the ion O22-. Oxygen compounds mainly have an oxidation state of -2.
Basic Features
Finally, we list the main properties of oxygen:
- Electron configuration: 1s22s22p4.
- Atomic number: 8.
- Atomic mass: 15.9994.
- Boiling point: -183.0 °C.
- Melting point: -218.4 °C.
- Density (if oxygen pressure is 1 atm at 0 °C): 1.429 g/l.
- Oxidation states: -1, -2, +2 (in compounds with fluorine).
