Oxygen

Properties of Oxygen

Oxygen is a highly reactive nonmetallic element; it is a strong oxidizing agent with high electronegativity and forms O2 at Standard Temperature and Pressure (STP).

Learning Objectives

Review the properties of oxygen.

Key Takeaways

Key Points

  • At standard temperature and pressure (STP), two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2.
  • Oxygen is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element. As such, it readily forms compounds (notably oxides) with almost all other elements.
  • Oxygen is a strong oxidizing agent and has the second-highest electronegativity of all reactive elements, second only to fluorine.
  • The solubility of oxygen in water is temperature-dependent; it condenses at 90.20 K and freezes at 54.36 K.

Key Terms

  • oxygen: A chemical element (symbol O) with an atomic number of 8 and atomic mass of 15.9994 amu.
  • ozone: An allotrope of oxygen (symbol O3) having three atoms in the molecule instead of the usual two; it is a blue gas, generated from oxygen by electrical discharge; it is poisonous and highly reactive, but it protects life on Earth by absorbing solar ultraviolet radiation in the upper atmosphere.
  • paramagnetic: Exhibiting paramagnetism (the tendency of magnetic dipoles to align with an external magnetic field).

Oxygen is an important part of the atmosphere and is necessary to sustain terrestrial life. Because it comprises most of the mass in water, it also comprises most of the mass of living organisms. All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Elemental oxygen (O2) is produced by cyanobacteria, algae, and plants through the process of photosynthesis, and is used in cellular respiration by most living organisms on earth. Oxygen is toxic to obligate anaerobic organisms (organisms which need a lack of oxygen for survival), which were the dominant form of early life on Earth, until O2 began to accumulate in the atmosphere.

Chemical Properties of Oxygen

At standard temperature and pressure (STP), two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element. As such, it readily forms compounds (notably, oxides) with almost all other elements. Oxygen is a strong oxidizing agent and has the second-highest electronegativity of all reactive elements, second only to fluorine. By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium, and the most abundant element by mass in the Earth’s crust, making up almost half of the crust’s mass. Free oxygen is too chemically reactive to appear on Earth without the photosynthetic action of living organisms, which use the energy of sunlight to produce elemental oxygen from water. Elemental O2 only began to accumulate in the atmosphere after the evolutionary appearance of photosynthetic organisms, roughly 2.5 billion years ago. Diatomic oxygen gas currently constitutes 20.8 percent of the volume of air.

Diatomic Oxygen

The two oxygen atoms in diatomic oxygen are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two and is often simplified in descriptions as a double bond, or as a combination of one two-electron bond and two three-electron bonds. Triplet oxygen (not to be confused with ozone, O3) is the ground state of the O2 molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond, in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.

In normal triplet form, O2 molecules are paramagnetic. This means they behave as magnets in the presence of an external magnetic field, because of the spin magnetic moments of the unpaired electrons in the molecule. Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet. Singlet oxygen is a name given to several higher-energy species of molecular O2 in which all the electron spins are paired. It is much more reactive toward common organic molecules than is the triplet form of molecular oxygen.

Physical Properties of Oxygen

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Liquid Oxygen: Oxygen bubbles rise through pale-blue liquid oxygen.

Oxygen is more soluble in water than nitrogen is; water contains approximately one molecule of O2 for every two molecules of N2, compared to an atmospheric ratio of approximately one to four. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg/L) dissolves at 0 °C than at 20 °C (7.6 mg/L). At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter. At 5 °C the solubility increases to 9.0 mL (50 percent more than at 25 °C) per liter for water and 7.2 mL (45 percent more) per liter for sea water.

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F). Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light).

High-purity liquid O2 is usually obtained by the fractional distillation of liquefied air. Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly reactive substance and must be segregated from combustible materials.

Oxides

An oxide is a chemical compound that contains at least one oxygen atom and one other element in its chemical formula.

Learning Objectives

Discuss the chemical properties of oxides.

Key Takeaways

Key Points

  • Metal oxides typically contain an anion of oxygen in the oxidation state of −2.
  • Noble metals (such as gold or platinum) are prized because they resist direct chemical combination with oxygen, and substances like gold (III) oxide must be generated by indirect routes.
  • The surface of most metals consists of oxides and hydroxides in the presence of air.
  • Metals tend to form basic oxides, non-metals tend to form acidic oxides, and amphoteric oxides are formed by elements near the boundary between metals and non-metals (metalloids).

Key Terms

  • passivation: The spontaneous formation of a hard non-reactive surface film (usually an oxide or nitride) that inhibits further corrosion.
  • oxide: A binary chemical compound of oxygen with another chemical element.
  • coke: Solid residue from roasting coal in a coke oven; used principally as a fuel and in the production of steel, and formerly as a domestic fuel.

Chemical Properties of Oxides

An oxide is a chemical compound that contains at least one oxygen atom and one other element in its chemical formula. Metal oxides typically contain an anion of oxygen in the oxidation state of −2. Most of the Earth’s crust consists of solid oxides, the result of elements being oxidized by the oxygen in air or water. Hydrocarbon combustion produces the two principal carbon oxides: carbon monoxide (CO) and carbon dioxide (CO2). Even materials considered pure elements often develop an oxide coating. For example, aluminum foil develops a thin skin of Al2O3 (called a passivation layer) that protects the foil from further corrosion.

Oxygen Exhibits High Reactivity

Due to its electronegativity, oxygen forms stable chemical bonds with almost all elements to give the corresponding oxides. Noble metals (such as gold or platinum) are prized because they resist direct chemical combination with oxygen, and substances like gold (III) oxide must be generated by indirect routes. Two independent pathways for corrosion of elements are hydrolysis and oxidation by oxygen. The combination of water and oxygen is even more corrosive. Virtually all elements burn in an atmosphere of oxygen or an oxygen-rich environment. In the presence of water and oxygen (or simply air), some elements—for example, sodium—react rapidly, even dangerously, to give hydroxide products. In part for this reason, alkali and alkaline earth metals are not found in nature in their metallic form. Cesium is so reactive with oxygen that it is used as a getter in vacuum tubes. Solutions of potassium and sodium, are used to deoxygenate and dehydrate some organic solvents.

Passivation

The surface of most metals consists of oxides and hydroxides in the presence of air. As mentioned above, a well-known example is aluminum foil, which is coated with a thin film of aluminium oxide that passivates the metal, slowing further corrosion. The aluminium oxide layer can be built to greater thickness by the process of electrolytic anodising. Though solid magnesium and aluminium react slowly with oxygen at STP, they, like most metals, burn in air, generating very high temperatures.

Polymeric vs. Monomeric Molecular Structures

Oxides of most metals adopt polymeric structures with M-O-M crosslinks. Because these crosslinks are strong, the solids tend to be insoluble in solvents, though they are attacked by acids and bases. The formulas are often deceptively simple. Many are nonstoichiometric compounds. In these oxides, the coordination number of the oxide ligand is 2 for most electronegative elements, and 3–6 for most metals.

Silicon Dioxide: Silicon dioxide (SiO2) is one of the most common oxides on the surface of earth. Like most oxides, it adopts a polymeric structure.

Although most metal oxides are polymeric, some oxides are monomeric molecules. The most famous molecular oxides are carbon dioxide and carbon monoxide. Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the formula being P4O10. Some polymeric oxides (selenium dioxide and sulfur trioxide) depolymerize to give molecules when heated. Tetroxides are rare, and there are only five known examples: ruthenium tetroxide, osmium tetroxide, hassium tetroxide, iridium tetroxide, and xenon tetroxide. Many oxyanions are known, such as polyphosphates and polyoxometalates. Oxycations are rarer, an example being nitrosonium (NO+). Of course many compounds are known with both oxides and other groups. For the transition metals, many oxo-complexes are known, as well as oxyhalides.

Acid-Base Reactions

Oxides can be attacked by acids and bases. Those attacked only by acids are basic oxides; those attacked only by bases are acidic oxides. Oxides that react with both acids and bases are amphoteric. Metals tend to form basic oxides, non-metals tend to form acidic oxides, and amphoteric oxides are formed by elements near the boundary between metals and non-metals (metalloids).

Other Redox Reactions

Metals are “won” from their oxides by chemical reduction. A common and cheap reducing agent is carbon in the form of coke. The most prominent example is that of iron ore smelting.

Oxides, such as iron (III) oxide (or rust, which consists of hydrated iron (III) oxides Fe2O3·nH2O and iron (III) oxide-hydroxide FeO(OH), Fe(OH)3), form when oxygen combines with iron.

Metal oxides can be reduced by organic compounds. This redox process is the basis for many important transformations in chemistry, such as the detoxification of drugs by the P450 enzymes and the production of ethylene oxide, which is converted to antifreeze. In such systems the metal center transfers an oxide ligand to the organic compound, followed by the regeneration of the metal oxide, often by oxygen in air.

Uses of Oxygen

Oxygen is essential for all aerobic organisms; common medical uses include oxygen therapy, hyperbaric medicine, and space suits.

Learning Objectives

Give examples of some common medically-related uses of oxygen.

Key Takeaways

Key Points

  • Oxygen is used in mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation.
  • Reactive oxygen species, such as superoxide ion (O2) and hydrogen peroxide (H2O2), are dangerous by-products of oxygen use in organisms.
  • Oxygen therapy not only increases oxygen levels in the patient’s blood, but also decreases resistance to blood flow in many types of diseased lungs, easing the work load on the heart. It is used to treat emphysema, pneumonia, and certain heart disorders (congestive heart failure).
  • Hyperbaric (high- pressure ) medicine uses special oxygen chambers to increase the partial pressure of O2 around the patient.

Key Terms

  • Hyperbaric: Of, relating to, or utilizing greater than normal pressure (as of oxygen).
  • aerobic: Organisms living in the presence of oxygen (e.g. aerobic bacteria).

Molecular dioxygen, O2, is essential for cellular respiration in all aerobic organisms. Oxygen is used in mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation. Reactive oxygen species, such as superoxide ion (O2) and hydrogen peroxide (H2O2), are dangerous by-products of oxygen use in organisms. Parts of the immune system of higher organisms, however, use reactive peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the hypersensitive response of plants against pathogen attack. An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute. This amounts to more than 6 billion tons of oxygen inhaled by humanity per year.

Oxygen Therapy

One of the medical uses of oxygen is oxygen therapy. Uptake of O2 from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Treatment not only increases oxygen levels in the patient’s blood, but has the secondary effect of decreasing resistance to blood flow in many types of diseased lungs, easing the work load on the heart. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders (congestive heart failure), some disorders that cause increased pulmonary artery pressure, as well as any disease that impairs the body’s ability to take up and use gaseous oxygen. Treatments are flexible enough to be used in hospitals, the patient’s home, or increasingly by portable devices. Oxygen tents were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of oxygen masks or nasal cannulas.

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Oxygen Concentrator: An oxygen concentrator for an emphysema patient.

Hyperbaric Medicine

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of O2 around the patient and, when needed, the medical staff. Carbon monoxide poisoning, gas gangrene, and decompression sickness (the ‘bends’) are sometimes treated using these devices. Increased O2 concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin. Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them. Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in their blood. Increasing the pressure of O2 as soon as possible is part of the treatment. Oxygen is also used medically for patients who require mechanical ventilation, often at concentrations above the 21% found in ambient air.

Use in Space Suits and Scuba Diving Suits

A notable application of O2 as a low-pressure breathing gas is in modern space suits, which surround their occupant’s body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressure of O2. This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits.

Scuba divers and submariners also rely on artificially delivered O2, but most often use normal pressure and/or mixtures of oxygen and air. O2 use in diving at higher than sea-level pressures is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~6 meters depth or less). Deeper diving requires significant dilution of O2 with other gases, such as nitrogen or helium, to help prevent oxygen toxicity. People who climb mountains or fly in non-pressurized fixed-wing aircrafts sometimes have supplemental O2 supplies.

Pressurized Commercial Airplanes

Passengers traveling in pressurized commercial airplanes have an emergency supply of O2 automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop. Pulling on the masks “to start the flow of oxygen,” as cabin safety instructions dictate, forces iron filings into a sample of sodium chlorate inside the canister where the reaction occurs. A steady stream of oxygen gas is then produced by the exothermic reaction.

Ozone

Ozone (O3) is diamagnetic (its electrons are all paired) and is a powerful oxidant.

Learning Objectives

Discuss the properties of ozone.

Key Takeaways

Key Points

  • Ozone is formed from O2 by the action of ultraviolet light and also atmospheric electrical discharges. It is present in low concentrations throughout the Earth’s atmosphere.
  • Ozone is slightly soluble in water, and much more soluble in inert nonpolar solvents such as carbon tetrachloride (CCl4) or fluorocarbons, where it forms a blue solution.
  • Ozone will oxidize most metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state.
  • Alkenes can be oxidatively cleaved by ozone, in a process called ozonolysis. With reductive workup (e.g., zinc in acetic acid or dimethyl sulfide), ketones and aldehydes will be formed. With oxidative workup (e.g. aqueous or alcoholic hydrogen peroxide), carboxylic acids will be formed.
  • Ozone, along with reactive forms of oxygen such as superoxide, singlet oxygen, hydrogen peroxide, and hypochlorite ions, is naturally produced by white blood cells and other biological systems as a means of destroying foreign bodies.

Key Terms

  • ozone: A triatomic molecule, also called trioxygen, consisting of three oxygen atoms (O3).
  • Alkenes: In organic chemistry, an alkene, olefin, is an unsaturated chemical compound containing at least one carbon-to-carbon double bond.
  • diamagnetic: Exhibiting diamagnetism; repelled by a magnet.

Properties of Ozone

Ozone (O3), or trioxygen, is a triatomic molecule consisting of three oxygen atoms. It is an allotrope of oxygen that is much less stable than the diatomic allotrope (O2), breaking down with a half life of about half an hour in the lower atmosphere to O2. Ozone is diamagnetic, which means that its electrons are all paired. In contrast, O2 is paramagnetic, containing two unpaired electrons.

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Resonance Structures of Ozone: The two resonance structures of O3 are shown.

Ozone in the Atmosphere

Ozone is formed from dioxygen by the action of ultraviolet light and also atmospheric electrical discharges. It is present in low concentrations throughout the Earth’s atmosphere. In total, ozone makes up only 0.6 parts per million of the atmosphere. Ozone’s odor is sharp, reminiscent of chlorine, and detectable by many people at concentrations of as little as 10 parts per billion in air. In standard conditions, ozone is a pale blue gas that condenses at progressively cryogenic temperatures to a dark blue liquid and finally a violet-black solid. Ozone is a powerful oxidant (far more so than dioxygen) and has many industrial and consumer applications related to oxidation. However, this same high oxidizing potential causes ozone to damage mucus and respiratory tissues in animals as well as tissues in plants, when it exists in concentrations above 100 parts per billion. This makes ozone a potent respiratory hazard and pollutant near ground level. However, the so-called ozone layer (a portion of the stratosphere with a higher concentration of ozone, from two to eight ppm) is beneficial. It prevents damaging ultraviolet light from reaching the Earth’s surface, which benefits all living organisms.

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Structure of Ozone: Ozone is a triatomic molecule with no unpaired electrons and a bent molecular shape. The bond lengths and angle formed by the three O atoms are shown.

Physical Properties of Ozone

Ozone is slightly soluble in water and much more soluble in inert nonpolar solvents such as carbon tetrachloride or fluorocarbons, where it forms a blue solution. At 161 K (−112 °C), it condenses to form a dark blue liquid. It is dangerous to allow this liquid to warm to its boiling point because both concentrated gaseous ozone and liquid ozone can detonate. At temperatures below 80 K (−193 °C), it forms a violet-black solid. It is also unstable at high concentrations, decaying to ordinary diatomic oxygen (with a half-life of about half an hour in atmospheric conditions):

[latex]2\text{O}_3 \rightarrow 3\text{O}_2[/latex]

This reaction proceeds more rapidly with increasing temperature and increased pressure.

Chemical Reactivity of Ozone

Ozone will oxidize most metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state. For example: [latex]{ 2\text{Cu} }^{ + }+2{ { \text{H} }_{ 3 }\text{O} }^{ + }+{ \text{O} }_{ 3 }\rightarrow 2{ \text{Cu} }^{ 2+ }+3{ \text{H} }_{ 2 }\text{O}+{ \text{O} }_{ 2 }[/latex]

Alkenes can be oxidatively cleaved by ozone in a process called ozonolysis, giving alcohols, aldehydes, ketones, and carboxylic acids, depending on the second step of the workup.

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Ozonolysis: The cleavage of carbon-carbon double bonds by O3 is shown in this figure.

Usually, ozonolysis is carried out in a solution of dichloromethane at a temperature of -78oC. After a sequence of cleavage and rearrangement, an organic ozonide is formed. With reductive workup (e.g., zinc in acetic acid or dimethyl sulfide), ketones and aldehydes will be formed. With oxidative workup (e.g., aqueous or alcoholic hydrogen peroxide), carboxylic acids will be formed.

Ozone’s Role in Biological Processes

Ozone, along with reactive forms of oxygen such as superoxide, singlet oxygen, hydrogen peroxide, and hypochlorite ions, is naturally produced by white blood cells and other biological systems (such as the roots of marigolds) as a means of destroying foreign bodies. Ozone reacts directly with organic double bonds.

When ozone breaks down to dioxygen, it produces oxygen free radicals, which are highly reactive and capable of damaging many organic molecules. Moreover, it is believed that the powerful oxidizing properties of ozone may be a contributing factor of inflammation. The cause-and-effect relationship of how the ozone is created in the body and what it does is still under consideration and still subject to various interpretations, since other body chemical processes can trigger some of the same reactions.