Oxygen
File:OxygenPoster.jpg Oxygen poster
Directed by	Richard Shepard
Produced by	Richard Shepard, Jonathan Stern, Mike Curb, Carole Curb Nemoy
Written by	Richard Shepard
Starring	Maura Tierney, Adrien Brody, Terry Kinney, Dylan Baker
Music by	Rolfe Kent
Distributed by	A-Pix Entertainment Inc.
Release date(s)	November 12, 1999
Running time	92 min.
Country	United States
Language	English

Oxygen is 1999 film, directed and written by Richard Shepard. The film follows a troubled and (unbeknownst to her husband) masochistic cop, Madeline Foster (played by Maura Tierney) as she pursues a kidnapper who calls himself Harry Houdini (Adrien Brody). The film bears a resemblance to the 1958 film Macabre.

Plot[link]

"Harry Houdini" (Adrien Brody) has kidnapped and buried alive a rich businessman's wife. He demands $1 million in cash from the woman's husband in return for the release of the location of her burial site before she runs out of oxygen. He also states that if police involvement is initiated, he will ensure she is never found alive. As expected, however, police involvement is initiated and Detective Madeline Foster (Maura Tierney) is assigned with the task of finding the buried woman as soon as possible, together with catching Houdini.

It is quickly realized though that Houdini, in fact, wanted the police to get involved so he could commit a second kidnapping - that of Foster. When Houdini is finally caught as the result of a car chase involving him and Detective Foster, he embroils her into his terrifying psychological world as she attempts to get him to reveal the location of his buried victim. When Foster's initial interrogation of Houdini proves unsuccessful, the FBI are assigned to take over the case, although they do not have any more luck. In a plot twist, it is revealed that Houdini has already committed a murder, that of his accomplice. He now has nothing to lose and the police realize the death of the buried woman would not be as significant as first thought.

However, Houdini offers them a final lifeline - if he is allowed to talk to Detective Foster, alone, he will allow his victim to be recovered alive. The police, now just puppets in Houdini's plan, give into his request and Foster is again embroiled in a psychological adventure of self-discovery with Houdini at the helm. Troubled, Foster begins to reveal her darker side to Houdini and he sees a kindred spirit in her. Just when she thinks she is getting through to him, Houdini persuades her that they are more alike than she ever imagined.

Soon after, it is revealed that Houdini was in fact a painter in the police department, and outlined the entire building in a very detailed map. Just as Foster is revealing the truth, Houdini takes her to the place where he buried the first woman. Foster escapes her hold, and releases the woman and takes Houdini's gun, then pushes him inside the coffin. His fate ends with him taunting Foster, so she shoots him rather than burying him.

The end of the movie is a short scene with Foster and her husband in a bar.

The film was shot on location in New York City.

External links[link]

• Oxygen at the Internet Movie Database

nitrogen ← oxygen → fluorine - ↑ O ↓ S 8O Periodic table
Appearance
Colorless gas; pale blue liquid. Oxygen bubbles rise in this photo of liquid oxygen. Spectral lines of oxygen
General properties
Name, symbol, number	oxygen, O, 8
Pronunciation	/ˈɒksɨdʒɨn/ OK-si-jin
Element category	nonmetal, chalcogen
Group, period, block	16, 2, p
Standard atomic weight	15.9994(3)
Electron configuration	1s2 2s2 2p4
Electrons per shell	2, 6 (Image)
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.429 g/L
Liquid density at b.p.	1.141 g·cm−3
Melting point	54.36 K, -218.79 °C, -361.82 °F
Boiling point	90.20 K, -182.95 °C, -297.31 °F
Critical point	154.59 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ·mol−1
Heat of vaporization	(O2) 6.82 kJ·mol−1
Molar heat capacity	(O2) 29.378 J·mol−1·K−1
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k at T (K)       61 73 90
Atomic properties
Oxidation states	2, 1, −1, −2
Electronegativity	3.44 (Pauling scale)
Ionization energies (more)	1st: 1313.9 kJ·mol−1
	2nd: 3388.3 kJ·mol−1
	3rd: 5300.5 kJ·mol−1
Covalent radius	66±2 pm
Van der Waals radius	152 pm
Miscellanea
Crystal structure	cubic
Magnetic ordering	paramagnetic
Thermal conductivity	26.58x10-3  W·m−1·K−1
Speed of sound	(gas, 27 °C) 330 m·s−1
CAS registry number	7782-44-7
Most stable isotopes
Main article: Isotopes of oxygen
iso NA half-life DM DE (MeV) DP 16O 99.76% 16O is stable with 8 neutrons 17O 0.039% 17O is stable with 9 neutrons 18O 0.201% 18O is stable with 10 neutrons
v t e · r

Oxygen ( /ˈɒksɨdʒɨn/ OK-si-jin) is the element with atomic number 8 and represented by the symbol O. Its name derives from the Greek roots ὀξύς (oxys) ("acid", literally "sharp", referring to the sour taste of acids) and -γενής (-genēs) ("producer", literally "begetter"), because at the time of naming, it was mistakenly thought that all acids required oxygen in their composition. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a very pale blue, 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 that readily forms compounds (notably oxides) with almost all other elements. Oxygen is a strong oxidizing agent and has the second-highest electronegativity of all the elements (only fluorine has a higher electronegativity).^[1] By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium^[2] and the most abundant element by mass in the Earth's crust, making up almost half of the crust's mass.^[3] 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 these organisms, roughly 2.5 billion years ago.^[4] Diatomic oxygen gas constitutes 20.8% of the volume of air.^[5]

Because it comprises most of the mass in water, oxygen comprises most of the mass of living organisms (for example, about two-thirds of the human body's mass^{[citation needed]}). 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 is produced by cyanobacteria, algae and plants, and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O₂ began to accumulate in the atmosphere. Another form (allotrope) of oxygen, ozone (O₃), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes atomic oxygen is a significant presence and a cause of erosion for spacecraft.^[6]

Oxygen was independently discovered by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his work was published first. The name oxygen was coined in 1777 by Antoine Lavoisier,^[7] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites with pressure-cycling to concentrate oxygen from air, electrolysis of water and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; oxygen therapy; and life support in aircraft, submarines, spaceflight and diving.

Characteristics[link]

Structure[link]

Oxygen discharge (spectrum) tube

At standard temperature and pressure, oxygen is a very pale blue, odorless gas with the molecular formula O₂, in which the two oxygen atoms 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 description as a double bond^[8] or as a combination of one two-electron bond and two three-electron bonds.^[9]

Triplet oxygen (not to be confused with ozone, O₃) is the ground state of the O₂ molecule.^[10] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.^[11] 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.^[10]

A trickle of liquid oxygen is deflected by a magnetic field, illustrating its paramagnetic property

In normal triplet form, O₂ molecules are paramagnetic. That is, they form a magnet in the presence of a magnetic field—because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O₂ molecules.^[12] 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.^[13][14]

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 towards common organic molecules than is molecular oxygen per se. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.^[15] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,^[16] and by the immune system as a source of active oxygen.^[17] Carotenoids in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.^[18]

Allotropes[link]

Ozone is a rare gas on Earth found mostly in the stratosphere.

The common allotrope of elemental oxygen on Earth is called dioxygen, O₂. It has a bond length of 121 pm and a bond energy of 498 kJ·mol⁻¹.^[19] This is the form that is used by complex forms of life, such as animals, in cellular respiration (see Biological role) and is the form that is a major part of the Earth's atmosphere (see Occurrence). Other aspects of O₂ are covered in the remainder of this article.

Trioxygen (O₃) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.^[20] Ozone is produced in the upper atmosphere when O₂ combines with atomic oxygen made by the splitting of O₂ by ultraviolet (UV) radiation.^[7] Since ozone absorbs strongly in the UV region of the spectrum, the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet.^[7] Near the Earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.^[20] The metastable molecule tetraoxygen (O₄) was discovered in 2001,^[21][22] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that this phase, created by pressurizing O₂ to 20 GPa, is in fact a rhombohedral O₈ cluster.^[23] This cluster has the potential to be a much more powerful oxidizer than either O₂ or O₃ and may therefore be used in rocket fuel.^[21][22] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa^[24] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.^[25]

Physical properties[link]

Oxygen is more soluble in water than nitrogen is; water contains approximately 1 molecule of O₂ for every 2 molecules of N₂, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L⁻¹) dissolves at 0 °C than at 20 °C (7.6 mg·L⁻¹).^[26][27] 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.^[28] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% 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).^[29] 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.^[30] 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.^[31]

Isotopes and stellar origin[link]

Late in a massive star's life, ¹⁶O concentrates in the O-shell, ¹⁷O in the H-shell and ¹⁸O in the He-shell.

Naturally occurring oxygen is composed of three stable isotopes, ¹⁶O, ¹⁷O, and ¹⁸O, with ¹⁶O being the most abundant (99.762% natural abundance).^[32]

Most ¹⁶O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process.^{[33] 17}O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.^[33] Most ¹⁸O is produced when ¹⁴N (made abundant from CNO burning) captures a ⁴He nucleus, making ¹⁸O common in the helium-rich zones of evolved, massive stars.^[33]

Fourteen radioisotopes have been characterized. The most stable are ¹⁵O with a half-life of 122.24 seconds and ¹⁴O with a half-life of 70.606 seconds.^[32] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds.^[32] The most common decay mode of the isotopes lighter than ¹⁶O is β⁺ decay^[34][35][36] to yield nitrogen, and the most common mode for the isotopes heavier than ¹⁸O is beta decay to yield fluorine.^[32]

Occurrence[link]

Oxygen is the most abundant chemical element, by mass, in the Earth's biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.^[2] About 0.9% of the Sun's mass is oxygen.^[5] Oxygen constitutes 49.2% of the Earth's crust by mass^[3] and is the major component of the world's oceans (88.8% by mass).^[5] Oxygen gas is the second most common component of the Earth's atmosphere, taking up 20.8% of its volume and 23.1% of its mass (some 10¹⁵ tonnes).^[5][38][39] Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere: Mars (with 0.1% O₂ by volume) and Venus have far lower concentrations. However, the O₂ surrounding these other planets is produced solely by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.

The unusually high concentration of oxygen gas on Earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, while respiration and decay remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.

Cold water holds more dissolved O₂.

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O₂ at lower temperatures (see Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.^[40] Polluted water may have reduced amounts of O₂ in it, depleted by decaying algae and other biomaterials through a process called eutrophication. Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O₂ needed to restore it to a normal concentration.^[41]

Biological role[link]

Photosynthesis and respiration[link]

Photosynthesis splits water to liberate O₂ and fixes CO₂ into sugar.

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. According to some estimates, Green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth and the rest is produced by terrestrial plants.^[42] Other estimates of the oceanic contribution to atmospheric oxygen are higher, while some estimates are lower, suggesting oceans produce ~45% of Earth's atmospheric oxygen each year.^[43]

A simplified overall formula for photosynthesis is:^[44]

6 CO₂ + 6 H₂O + photons → C₆H₁₂O₆ + 6 O₂ (or simply carbon dioxide + water + sunlight → glucose + dioxygen)

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.^[45] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize ATP via photophosphorylation.^[46] The O₂ remaining after oxidation of the water molecule is released into the atmosphere.^[47]

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. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:

C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O + 2880 kJ·mol⁻¹

In vertebrates, O₂ diffuses through membranes in the lungs and into red blood cells. Hemoglobin binds O₂, changing its color from bluish red to bright red^[20] (CO₂ is released from another part of hemoglobin through the Bohr effect). Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).^[38] A liter of blood can dissolve 200 cm³ of O₂.^[38]

Reactive oxygen species, such as superoxide ion (O−
2) and hydrogen peroxide (H₂O₂), are dangerous by-products of oxygen use in organisms.^[38] Parts of the immune system of higher organisms, however, create 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.^[46]

An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute.^[48] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.^[49]

Content in body[link]

The oxygen content in the body of a living organism is usually highest in the respiratory system, and decreases along any arterial system, peripheral tissues and venous system, respectively. Oxygen content in this sense is often given as the partial pressure, which is the pressure which oxygen would have if it alone occupied the volume.^[53]

Build-up in the atmosphere[link]

O₂ build-up in Earth's atmosphere: 1) no O₂ produced; 2) O₂ produced, but absorbed in oceans & seabed rock; 3) O₂ starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5) O₂ sinks filled and the gas accumulates

Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved. Free oxygen first appeared in significant quantities during the Paleoproterozoic eon (between 2.5 and 1.6 billion years ago). At first, the oxygen combined with dissolved iron in the oceans to form banded iron formations. Free oxygen started to outgas from the oceans 2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.^[54]

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the Great Oxygenation Event(oxygen catastrophe) about 2.4 billion years ago. However, cellular respiration using O₂ enables aerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's biosphere.^[55] Photosynthesis and cellular respiration of O₂ allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian period 540 million years ago, O₂ levels have fluctuated between 15% and 30% by volume.^[56] Towards the end of the Carboniferous period (about 300 million years ago) atmospheric O₂ levels reached a maximum of 35% by volume,^[56] which may have contributed to the large size of insects and amphibians at this time.^[57] Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.^[12] At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O₂ in the present atmosphere.^[58]

History[link]

Early experiments[link]

Philo's experiment inspired later investigators.

One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.^[59] Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration.^[60]

In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow (1641–1679) refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.^[61] In one experiment he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.^[62] From this he surmised that nitroaereus is consumed in both respiration and combustion.

Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it.^[61] He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.^[61] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".^[62]

Phlogiston theory[link]

Stahl helped develop and popularize the phlogiston theory.

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as an element.^[26] This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes.

Established in 1667 by the German alchemist J. J. Becher, and modified by the chemist Georg Ernst Stahl by 1731,^[63] phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx.^[60]

Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; whereas non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.^[60] The fact that a substance like wood actually gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of the first clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they were supposedly losing phlogiston).

Discovery[link]

Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.

Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various nitrates by about 1772.^[5][60] Scheele called the gas "fire air" because it was the only known supporter of combustion, and wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.^[64]

Joseph Priestley is usually given priority in the discovery.

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named "dephlogisticated air".^[5] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."^[26] Priestley published his findings in 1775 in a paper titled "An Account of Further Discoveries in Air" which was included in the second volume of his book titled Experiments and Observations on Different Kinds of Air.^[60][65] Because he published his findings first, Priestley is usually given priority in the discovery.

The noted French chemist Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30, 1774 that described his own discovery of the previously unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).^[64]