oxygen

 Carl Wilhelm Scheele, a Swedish scientist, discovered oxygen in 1772. He did it by heating potassium nitrate, mercuric oxide, and a variety of other materials. Three years prior to Scheele's publication, in 1774, English scientist Joseph Priestley independently found oxygen by the thermal breakdown of mercuric oxide. He reported his results in the same year.

The phlogiston theory, which had been accepted up until that point, was disproved by the French chemist Antoine-Laurent Lavoisier in 1775–1780. Lavoisier noted that oxygen had a tendency to form acids through its combination with a variety of substances, and as a result, he named the element oxygen (oxygène), derived from the Greek words for "acid former."

Occurrence and properties:

                                The most common element in the crust of the Earth is oxygen, which makes up 46% of its mass. In saltwater, the percentage of oxygen by weight is 89%, while it is 21% by volume in the atmosphere. It is combined with metals and nonmetals in rocks to form acidic (sulfur, carbon, aluminum, and phosphorus) or basic (calcium, magnesium, and iron) oxides as well as salt-like compounds (sulfates, carbonates, silicates, aluminates, and phosphates), which can be thought of as derived from the acidic and basic oxides. Despite their abundance, these solid complexes cannot be utilized as oxygen sources due to the high cost involved in separating the element from their tightly bound metal atom combinations.

Oxygen is a pale blue liquid at −183 °C (−297 °F); it solidifies around roughly −218 °C (−361 °F). Air is 1.1 times lighter than pure oxygen.

While certain microorganisms and animals absorb oxygen from the environment during respiration and return it to it as carbon dioxide, green plants use photosynthesis to digest carbon dioxide in the presence of sunshine and produce free oxygen. Photosynthesis is responsible for almost all of the atmospheric free oxygen. In 100 parts fresh water at 20 °C (68 °F), around 3 parts oxygen by volume dissolves; somewhat less in seawater. Fish and other marine creatures depend on dissolved oxygen to breathe.

The stable isotopes of oxygen-16 (99.759 percent), oxygen-17 (0.037 percent), and oxygen-18 (0.204 percent) are combined to form atmospheric oxygen. There are known to be several radioactive isotopes created artificially. The longest-lived, oxygen-15, has been employed to research animal respiration (half-life of 124 seconds).

Allotropy:

Diatomic (O2) and triatomic (O3, ozone) are the two allotropic forms of oxygen. The diatomic form's characteristics imply that oxygen's paramagnetism is caused by six electrons joining the atoms and two electrons remaining unpaired. The ozone molecule's three atoms are not arranged in a straight line.

The following equation indicates how oxygen may be converted into ozone:

The process as described is endothermic, meaning energy must be supplied in order for it to continue; the presence of transition metals or their oxides encourages the conversion of ozone back into diatomic oxygen. A silent electrical discharge converts pure oxygen partially into ozone.

 The reaction is also triggered by the absorption of ultraviolet light with wavelengths of about 250 nanometers (nm, or 10−9 meters). This process occurs in the upper atmosphere and eliminates radiation that would be dangerous for life on Earth's surface. The strong smell of ozone is detectable in small spaces where electrical equipment sparks, such as in generator rooms. Light blue in color, ozone has a density 1.658 times that of air.

Ozone is a potent oxidizing agent that may change a variety of organic molecules into oxygenated derivatives including aldehydes and acids, as well as sulfur dioxide to sulfur trioxide, sulfides to sulfates, and iodides to iodine (offering an analytical method for its measurement). Smog gets its unpleasant quality from the conversion of hydrocarbons from car exhaust fumes to these acids and aldehydes by ozone. Ozone has been utilized in the chemical reagent, disinfectant, sewage treatment, water purification, and textile bleaching industries on a commercial scale.

Preparative methods:

The quantity of oxygen that is desired determines the production methods that are selected. Among the processes used in laboratories are the following:

1. Thermal decomposition of certain salts, including potassium chlorate or potassium nitrate: Manganese dioxide (pyrolusite, MnO2) is a common oxide of transition metal that catalyzes the decomposition of potassium chlorate. The catalyst lowers the temperature required to affect the evolution of oxygen from 400 °C to 250 °C.

2. Thermal breakdown of heavy metal oxides: Schoele and Priestley employed mercury(II) oxide to prepare oxygen.

3. Thermal decomposition of metal peroxides or hydrogen peroxide: As the equations illustrate, an early commercial method for producing hydrogen peroxide or extracting oxygen from the environment relied on the creation of barium peroxide from the oxide.

4. To enable electric current conduction, electrolyze water that contains trace amounts of salts or acids:

Commercial production and use:

Liquid air is fractionally distilled to prepare oxygen when tonnage requirements are met. Since oxygen has a higher boiling point than nitrogen and argon, it is less volatile among the primary constituents of air. The procedure makes use of the fact that compressed gases cool when given the chance to expand.

 The following are important steps in the operation: The process involves filtering the air to remove particles, absorbing moisture and carbon dioxide in alkali, compressing the air and removing the heat generated by compression using standard cooling techniques, passing the compressed and cooled air through coils inside a chamber, and allowing a portion of the compressed air to expand at a pressure of approximately 200 atmospheres.in the chamber, cooling the coils;

 (6) the expanded gas is fed back into the compressor, where it undergoes several more stages of expansion and compression until the compressed air liquefies at a temperature of -196 °C;

 (7) the liquid air is heated to separate the liquid oxygen from the light, rare gases and then the nitrogen. The product obtained from many fractionations is sufficiently pure (99.5 percent) for the majority of industrial uses.

In "blowing" high carbon steel, or volatilizing carbon dioxide and other nonmetal impurities in a quicker and easier-to-control process than if air were used, the steel industry is the biggest user of pure oxygen. Compared to other chemical procedures, oxygen treatment of sewage shows potential for more effective treatment of liquid effluents. Waste incineration with pure oxygen in closed systems has gained importance. Liquid oxygen, or "LOX," is a component of rocket oxidizer fuels. How much LOX is used depends on space program activities. Diving bells and submarines both need pure oxygen.

In the chemical industry, commercial oxygen, or oxygen-enriched air, has taken the place of regular air in the production of oxidation-controlled compounds like methanol, ethylene oxide, and acetylene. Oxygen has medical uses such as inhalators, oxygen tents, and incubators for young patients. During general anesthesia, oxygen-enriched gaseous anesthetics guarantee life support. An important component of many kiln-using businesses is oxygen.

Chemical properties and reactions:

Oxygen's electronegativity and electron affinity are big values, which are typical of elements that exhibit only nonmetallic behavior. Because of the two partially full outer orbitals, oxygen naturally assumes a negative oxidation state in all of its compounds. The oxide ion O2− is produced when electron transport fills these orbitals. 

It is thought that every oxygen in peroxides (species that contain the ion O22−) has a charge of −1. An oxidizing agent's ability to receive electrons through full or partial transfer is what makes it such. Such an agent lowers its own oxidation state when it interacts with an electron-donating material. Reduction is the term used to describe the shift (lowering) of oxygen from its zero state to its -2 state. One may think of oxygen as thought of as the "original" oxidizing agent, with oxidation and reduction being named after its characteristic behavior.

Under typical circumstances, oxygen forms the triatomic species ozone (O3) and the diatomic species O2, as was discussed in the section on allotropy. Some evidence points to O4, a very unstable tetratomic species. There are two unpaired electrons in antibonding orbitals in the molecular diatomic form. The existence of these electrons is confirmed by oxygen's paramagnetic behavior.

One theory to explain the extreme reactivity of ozone is that one of the three oxygen atoms is in a condition known as "atomicity," which causes the atom to react and separate from the O3 molecule to leave molecular oxygen behind.

O2 is a molecular species that is not particularly thought of as the "original" oxidizing agent, with oxidation and reduction being named after its characteristic behavior.

Under typical circumstances, oxygen forms the triatomic species ozone (O3) and the diatomic species O2, as was discussed in the section on allotropy. Some evidence points to O4, a very unstable tetratomic species. There are two unpaired electrons in antibonding orbitals in the molecular diatomic form. The preactive under standard (ambient) pressure and temperature conditions. The atomic species O has a far higher reactivity. O2 → 2O has a significant dissociation energy of 117.2 kcal/mol.

Most of the substances that include oxygen have an oxidation state of −2. It creates a wide variety of compounds with covalent bonds, including recognized as the "original" oxidizing agent, and the terminology used to explain oxidation and reduction is predicated on this characteristic behavior of oxygen.

According to the explanation in the allotropy section, oxygen normally forms the triatomic species ozone (O3) and the diatomic species O2. A very unstable tetratomic species, O4, has been suggested by some evidence. Two unpaired electrons that are in antibonding orbitals are present in the molecule diatomic form. the preactive at standard (ambient) pressures and temperatures. Significantly more reactive is the atomic species O. At 117.2 kilocalories per mole, the dissociation energy (O2 → 2O) is high.

In the majority of its compounds, oxygen has an oxidation state of −2. It creates a wide variety of covalently bound molecules, including har oxides of nonmetals, such as water (H2O), sulfur dioxide (SO2), and carbon dioxide (CO2); organic compounds such as alcohols, aldehydes, and carboxylic acids; common acids such as sulfuric (H2SO4), carbonic (H2CO3), and nitric (HNO3); and corresponding salts, such as sodium sulfate (Na2SO4), sodium carbonate (Na2CO3), and sodium nitrate (NaNO3). Oxygen is present as the oxide ion, O2-, in the crystalline structure of solid metallic oxides such as calcium oxide, CaO. Metallic superoxides, such as potassium superoxide, KO2, contain the O2- ion, whereas metallic peroxides, such as barium peroxide, BaO2, contain the O22- ion.