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= Oxygen =
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Introduction
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Oxygen is a chemical element; it has symbol O and atomic number 8. It
is a member of the chalcogen group in the periodic table, a highly
reactive nonmetal, and a potent oxidizing agent that readily forms
oxides with most elements as well as with other compounds. Oxygen is
the most abundant element in Earth's crust, making up almost half of
the Earth's crust in the form of various oxides such as water, carbon
dioxide, iron oxides and silicates. It is the third-most abundant
element in the universe after hydrogen and helium.
At standard temperature and pressure, two oxygen atoms will bind
covalently to form dioxygen, a colorless and odorless diatomic gas
with the chemical formula . Dioxygen gas currently constitutes
approximately 20.95% molar fraction of the Earth's atmosphere, though
this has changed considerably over long periods of time in Earth's
history. A much rarer triatomic allotrope of oxygen, ozone (),
strongly absorbs the UVB and UVC wavelengths and forms a protective
ozone layer at the lower stratosphere, which shields the biosphere
from ionizing ultraviolet radiation. However, ozone present at the
surface is a corrosive byproduct of smog and thus an air pollutant.
All eukaryotic organisms, including plants, animals, fungi, algae and
most protists, need oxygen for cellular respiration, a process that
extracts chemical energy by the reaction of oxygen with organic
molecules derived from food and releases carbon dioxide as a waste
product.
Many major classes of organic molecules in living organisms contain
oxygen atoms, such as proteins, nucleic acids, carbohydrates and fats,
as do the major constituent inorganic compounds of animal shells,
teeth, and bone. Most of the mass of living organisms is oxygen as a
component of water, the major constituent of lifeforms. Oxygen in
Earth's atmosphere is produced by biotic photosynthesis, in which
photon energy in sunlight is captured by chlorophyll to split water
molecules and then react with carbon dioxide to produce carbohydrates
and oxygen is released as a byproduct. Oxygen is too chemically
reactive to remain a free element in air without being continuously
replenished by the photosynthetic activities of autotrophs such as
cyanobacteria, chloroplast-bearing algae and plants.
Oxygen was isolated by Michael Sendivogius before 1604, but it is
commonly believed that the element was discovered independently by
Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph
Priestley in Wiltshire, in 1774. Priority is often given for Priestley
because his work was published first. Priestley, however, called
oxygen "dephlogisticated air", and did not recognize it as a chemical
element. The name 'oxygen' was coined in 1777 by Antoine Lavoisier,
who first recognized oxygen as a chemical element and correctly
characterized the role it plays in combustion.
Common industrial uses of oxygen include production of steel, plastics
and textiles, brazing, welding and cutting of steels and other metals,
rocket propellant, oxygen therapy, and life support systems in
aircraft, submarines, spaceflight and diving.
History of study
======================================================================
The modern concept of the element oxygen developed over five centuries
and included many related discoveries and unsuccessful theories.
Multiple people made different contributions to the concept: no one
person discovered oxygen.
Early experiments
===================
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 ', 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. 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 Ibn al-Nafis, writing in 1250CE, correctly
described oxygenation of blood in the circulatory system; Michael
Servetus rediscovered this concept in 1553 but his books were
systematically destroyed. A scientifically based and influential
description was published by William Harvey in 1628.
Leonardo da Vinci observed that a portion of air is consumed during
combustion and respiration.
Polish alchemist, philosopher, and physician Michael Sendivogius
(Michał Sędziwój), writing in 1604, described a substance contained in
air, referring to it as ('food of life'); this substance is identical
with oxygen. During his experiments, performed between 1598 and 1604,
Sendivogius properly recognized that the substance is equivalent to
the gaseous byproduct released by the thermal decomposition of
potassium nitrate. However, this important connection was not
understood by contemporary scientists like Robert Boyle.
Unaware of Sendivogius's work, John Mayow wrote about a portion of air
that provided heat in a fire and the human body. This work was ignored
because it failed to align with the prevailing phlogiston theory of
air and fire.
Mayow observed that antimony increased in weight when heated, and
inferred that the must have combined with it. He also thought that
the lungs separate from air and pass it into the blood and that
animal heat and muscle movement result from the reaction of with
certain substances in the body. Accounts of these and other
experiments and ideas were published in 1668 in his work ' in the
tract "".
After Robert Boyle proved that air is necessary for combustion in the
late 17th century, English chemist John Mayow (1641-1679) refined this
work by showing that fire requires only a part of air that he called .
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. From this, he surmised that is consumed in both respiration
and combustion.
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 a chemical element. 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, 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.
Highly combustible materials that leave little residue, such as wood
or coal, were thought to be made mostly of phlogiston; 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.
Scientific era
================
Swedish pharmacist Carl Wilhelm Scheele produced and described some
properties of oxygen sometime around 1770-1775, but did not publish
his work until a few years later because he was unable to interpret
his work in the framework of the phlogiston theory. Scheele had
produced oxygen gas by heating mercuric oxide (HgO) and various
nitrates in 1771-1772. After reading about Priestley's work in 1775,
Scheele published in 1777, calling the gas "fire air" because it was
then the only known agent to support combustion.
In the meantime, on August 1, 1774, an experiment conducted by the
British clergyman Joseph Priestley focused sunlight on mercuric oxide
contained in a glass tube, which liberated a gas he named
"dephlogisticated air". 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, Priestley 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." 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'.
The French chemist Antoine Lavoisier later claimed to have discovered
the new substance independently. Priestley visited Lavoisier in
October 1774 and told him about his experiment and how he liberated
the new gas. Scheele had also dispatched a letter to Lavoisier on
September 30, 1774, which described his 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).
Discrediting Philogiston theory
=================================
Lavoisier conducted the first adequate quantitative experiments on
oxidation and gave the first correct explanation of how combustion
works. He used these and similar experiments, all started in 1774, to
discredit the phlogiston theory and to prove that the substance
discovered by Priestley and Scheele was a chemical element.
In one experiment, Lavoisier observed that there was no overall
increase in weight when tin and air were heated in a closed container.
He noted that air rushed in when he opened the container, which
indicated that part of the trapped air had been consumed. He also
noted that the tin had increased in weight and that increase was the
same as the weight of the air that rushed back in. This and other
experiments on combustion were documented in his book ', which was
published in 1777. In that work, he proved that air is a mixture of
two gases; 'vital air', which is essential to combustion and
respiration, and (from Greek 'lifeless'), which did not support
either. later became 'nitrogen' in English, although it has kept the
earlier name in French and several other European languages.
Etymology
===========
Lavoisier renamed "vital air" to in 1777 from the Greek roots (;
"acid", literally 'sharp', from the taste of acids) and (;
"producer", literally 'begetter'), because he mistakenly believed that
oxygen was a constituent of all acids. Chemists (such as Sir Humphry
Davy in 1812) eventually determined that Lavoisier was wrong in this
regard (e.g. Hydrogen chloride (HCl) is a strong acid that does not
contain oxygen), but by then the name was too well established.
'Oxygen' entered the English language despite opposition by English
scientists and the fact that the Englishman Priestley had first
isolated the gas and written about it. This is partly due to a poem
praising the gas titled "Oxygen" in the popular book 'The Botanic
Garden' (1791) by Erasmus Darwin, grandfather of Charles Darwin.
Later history
===============
John Dalton's original atomic hypothesis presumed that all elements
were monatomic and that the atoms in compounds would normally have the
simplest atomic ratios with respect to one another. For example,
Dalton assumed that water's formula was HO, leading to the conclusion
that the atomic mass of oxygen was 8 times that of hydrogen, instead
of the modern value of about 16. In 1805, Joseph Louis Gay-Lussac and
Alexander von Humboldt showed that water is formed of two volumes of
hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had
arrived at the correct interpretation of water's composition, based on
what is now called Avogadro's law and the diatomic elemental molecules
in those gases.
In 1879 the French brothers Quentin and Arthur
Brin discovered a commercially viable reaction to create oxygen. They
realized that the known reversible reaction
(s) + (g) ↔ (s) was deactivated by the formation of barium carbonate
from carbon dioxide in the air; treating air to remove the carbon
dioxide allowed the reaction be reversed indefinitely. Their company
used the process between 1886 and 1906 when more economical fractional
distillation began to be used.
By the late 19th century scientists realized that air could be
liquefied and its components isolated by compressing and cooling it.
Using a cascade method, Swiss chemist and physicist Raoul Pierre
Pictet evaporated liquid sulfur dioxide in order to liquefy carbon
dioxide, which in turn was evaporated to cool oxygen gas enough to
liquefy it. He sent a telegram on December 22, 1877, to the French
Academy of Sciences in Paris announcing his discovery of liquid
oxygen. Just two days later, French physicist Louis Paul Cailletet
announced his own method of liquefying molecular oxygen. Only a few
drops of the liquid were produced in each case and no meaningful
analysis could be conducted. Oxygen was liquefied in a stable state
for the first time on March 29, 1883, by Polish scientists from
Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.
In 1891 Scottish chemist James Dewar was able to produce enough liquid
oxygen for study. The first commercially viable process for producing
liquid oxygen was independently developed in 1895 by German engineer
Carl von Linde and British engineer William Hampson. Both men lowered
the temperature of air until it liquefied and then distilled the
component gases by boiling them off one at a time and capturing them
separately. Later, in 1901, oxyacetylene welding was demonstrated for
the first time by burning a mixture of acetylene and compressed . This
method of welding and cutting metal later became common.
In 1923, the American scientist Robert H. Goddard became the first
person to develop a rocket engine that burned liquid fuel; the engine
used gasoline for fuel and liquid oxygen as the oxidizer. Goddard
successfully flew a small liquid-fueled rocket 56 m at 97 km/h on
March 16, 1926, in Auburn, Massachusetts, US.
Properties and molecular structure
====================================
At standard temperature and pressure, oxygen is a colorless, odorless,
and tasteless gas with the molecular formula , referred to as
dioxygen.
As 'dioxygen', two oxygen atoms are chemically bound to each other.
The bond can be variously described based on level of theory, but is
reasonably and simply described as a covalent double bond that results
from the filling of molecular orbitals formed from the atomic orbitals
of the individual oxygen atoms, the filling of which results in a bond
order of two. More specifically, the double bond is the result of
sequential, low-to-high energy, or Aufbau, filling of orbitals, and
the resulting cancellation of contributions from the 2s electrons,
after sequential filling of the low σ and σ* orbitals; σ overlap of
the two atomic 2p orbitals that lie along the O-O molecular axis and π
overlap of two pairs of atomic 2p orbitals perpendicular to the O-O
molecular axis, and then cancellation of contributions from the
remaining two 2p electrons after their partial filling of the π*
orbitals.
This combination of cancellations and σ and π overlaps results in
dioxygen's double-bond character and reactivity, and a triplet
electronic ground state. An electron configuration with two unpaired
electrons, as is found in dioxygen orbitals (see the filled π*
orbitals in the diagram) that are of equal energy--i.e.,
degenerate--is a configuration termed a spin triplet state. Hence, the
ground state of the molecule is referred to as triplet oxygen. The
highest-energy, partially filled orbitals are antibonding, and so
their filling weakens the bond order from three to two. Because of its
unpaired electrons, triplet oxygen reacts only slowly with most
organic molecules, which have paired electron spins; this prevents
spontaneous combustion.
In the triplet form, molecules are paramagnetic. That is, they impart
magnetic character to oxygen when it is 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
molecules. Liquid oxygen is so magnetic that, in laboratory
demonstrations, a bridge of liquid oxygen may be supported against its
own weight between the poles of a powerful magnet. Oxygen's
paramagnetism can be used analytically in paramagnetic oxygen gas
analysers that determine gaseous oxygen concentration, especially in
industrial process control and medicine.
Singlet oxygen is a name given to several higher-energy species of
molecular in which all the electron spins are paired. It is much more
reactive with common organic molecules than is normal (triplet)
molecular oxygen. In nature, singlet oxygen is commonly formed from
water during photosynthesis, using the energy of sunlight. It is also
produced in the troposphere by the photolysis of ozone by light of
short wavelength and by the immune system as a source of active
oxygen. Carotenoids in photosynthetic organisms (and possibly 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.
Allotropes
============
The common allotrope of elemental oxygen on Earth is called dioxygen,
, the major part of the Earth's atmospheric oxygen (see Occurrence).
O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.
Trioxygen () is usually known as ozone and is a very reactive
allotrope of oxygen that is damaging to lung tissue. Ozone is produced
in the upper atmosphere when combines with atomic oxygen made by the
splitting of by ultraviolet (UV) radiation. 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. Near the Earth's surface, it is a pollutant formed as a
by-product of automobile exhaust. At low earth orbit altitudes,
sufficient atomic oxygen is present to cause corrosion of spacecraft.
The metastable molecule tetraoxygen () was discovered in 2001, and was
assumed to exist in one of the six phases of solid oxygen. In 2006
this phase, created by pressurizing to 20 GPa, was shown to form a
rhombohedral cluster. This cluster has the potential to be a much
more powerful oxidizer than either or and may therefore be used in
rocket fuel. A metallic phase was discovered in 1990 when solid oxygen
is subjected to a pressure of above 96 GPa and it was shown in 1998
that at very low temperatures, this phase becomes superconducting.
Physical properties
=====================
Oxygen dissolves more readily in water than nitrogen does. Water in
equilibrium with air contains approximately 1 molecule of dissolved
for every 2 molecules of (1:2), compared with an atmospheric ratio of
approximately 1:4. The solubility of oxygen in water is
temperature-dependent, and about twice as much () dissolves at than
at ().
At and 1 atm of air, freshwater can dissolve about 6.04 milliliters
(mL) of oxygen per liter, and seawater contains about 4.95 mL per
liter.
At the solubility increases to 9.0 mL (50% more than at ) per liter
for freshwater and 7.2 mL (45% more) per liter for sea water.
Oxygen gas dissolved in water at sea-level (milliliters per liter)
! ! !
|Freshwater |9.00 |6.04
|Seawater |7.20 |4.95
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 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 is usually obtained by
the fractional distillation of liquefied air. Liquid oxygen may also
be condensed from air using liquid nitrogen as a coolant. Liquid
oxygen is a highly reactive substance and must be segregated from
combustible materials.
The spectroscopy of molecular oxygen is associated with the
atmospheric processes of aurora and airglow. The absorption in the
Herzberg continuum and Schumann-Runge bands in the ultraviolet
produces atomic oxygen that is important in the chemistry of the
middle atmosphere. Excited-state singlet molecular oxygen is
responsible for red chemiluminescence in solution.
Table of thermal and physical properties of oxygen (O2) at atmospheric
pressure:
|Temperature (K) |Density (kg/m3) |Specific heat (kJ/(kg·K)) |Dynamic
viscosity (kg/(m·s)) |Kinematic viscosity (m2/s) |Thermal conductivity
(W/(m·K)) |Thermal diffusivity (m2/s) |Prandtl Number
|100 |3.945 |0.962 |7.64E-06 |1.94E-06 |0.00925 |2.44E-06 |0.796
|150 |2.585 |0.921 |1.15E-05 |4.44E-06 |0.0138 |5.80E-06 |0.766
|200 |1.93 |0.915 |1.48E-05 |7.64E-06 |0.0183 |1.04E-05 |0.737
|250 |1.542 |0.915 |1.79E-05 |1.16E-05 |0.0226 |1.60E-05 |0.723
|300 |1.284 |0.92 |2.07E-05 |1.61E-05 |0.0268 |2.27E-05 |0.711
|350 |1.1 |0.929 |2.34E-05 |2.12E-05 |0.0296 |2.90E-05 |0.733
|400 |0.962 |1.0408 |2.58E-05 |2.68E-05 |0.033 |3.64E-05 |0.737
|450 |0.8554 |0.956 |2.81E-05 |3.29E-05 |0.0363 |4.44E-05 |0.741
|500 |0.7698 |0.972 |3.03E-05 |3.94E-05 |0.0412 |5.51E-05 |0.716
|550 |0.6998 |0.988 |3.24E-05 |4.63E-05 |0.0441 |6.38E-05 |0.726
|600 |0.6414 |1.003 |3.44E-05 |5.36E-05 |0.0473 |7.35E-05 |0.729
|700 |0.5498 |1.031 |3.81E-05 |6.93E-05 |0.0528 |9.31E-05 |0.744
|800 |0.481 |1.054 |4.15E-05 |8.63E-05 |0.0589 |1.16E-04 |0.743
|900 |0.4275 |1.074 |4.47E-05 |1.05E-04 |0.0649 |1.41E-04 |0.74
|1000 |0.3848 |1.09 |4.77E-05 |1.24E-04 |0.071 |1.69E-04 |0.733
|1100 |0.3498 |1.103 |5.06E-05 |1.45E-04 |0.0758 |1.96E-04 |0.736
|1200 |0.3206 |1.0408 |5.33E-05 |1.661E-04 |0.0819 |2.29E-04 |0.725
|1300 |0.296 |1.125 |5.88E-05 |1.99E-04 |0.0871 |2.62E-04 |0.721
Isotopes and stellar origin
=============================
Naturally occurring oxygen is composed of three stable isotopes, 16O,
17O, and 18O, with 16O being the most abundant (99.762% natural
abundance).
16O is the one of the dominant fusion products in massive stars. It is
synthesized at the end of the helium fusion process with some
synthesis in the neon burning process. Both 17O and 18O require seed
nuclei. 17O 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. 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.
Fifteen radioisotopes have been characterized, ranging from 11O to
28O. The most stable are 15O with a half-life of 122.24 seconds and
14O with a half-life of 70.606 seconds. All of the remaining
radioactive isotopes have half-lives that are less than 27 seconds and
the majority of these have half-lives that are less than 83
milliseconds. The most common decay mode of the isotopes lighter than
16O is β+ decay to yield nitrogen, and the most common mode for the
isotopes heavier than 18O is beta decay to yield fluorine.
Occurrence
============
Ten most common elements in the Milky Way Galaxy estimated
spectroscopically
!Z !! Element !! colspan="2"|Mass fraction in parts per million
1 Hydrogen align="right"|
2 Helium align="right"|
8 Oxygen align="right"|
6 Carbon align="right"|
10 Neon align="right"|
26 Iron align="right"|
7 Nitrogen align="right"|
14 Silicon align="right"|
12 Magnesium align="right"|
16 Sulfur align="right"|
Oxygen is the third most abundant chemical element in the universe,
after hydrogen and helium. About 0.9% of the Sun's mass is oxygen.
Oxygen constitutes 49.2% of the Earth's crust by mass as part of oxide
compounds such as silicon dioxide and is the most abundant element by
mass in the Earth's crust. It is also the major component of the
world's oceans (88.8% by mass). 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).
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%
by volume) and Venus have much less. The surrounding those planets is
produced solely by the action of ultraviolet radiation on
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, decay, and combustion remove it from
the atmosphere. In the present equilibrium, production and consumption
occur at the same rate.
Oxygen levels in the atmosphere are trending slightly downward
globally, possibly because of fossil-fuel burning.
Free oxygen also occurs in solution in the world's water bodies. The
increased solubility of 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. Scientists assess this aspect of water quality by measuring
the water's biochemical oxygen demand, or the amount of needed to
restore it to a normal concentration.
Significant deoxygenation has been observed in tropical oceans.
Warming oceans waters are expected to lose oxygen over the next
century and into the future for a thousand years; the possible
consequences include minimal oxygen zones which are unable to support
macrofauna.
Analysis
==========
Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in
the shells and skeletons of marine organisms to determine the climate
millions of years ago (see oxygen isotope ratio cycle). Seawater
molecules that contain the lighter isotope, oxygen-16, evaporate at a
slightly faster rate than water molecules containing the 12% heavier
oxygen-18, and this disparity increases at lower temperatures. During
periods of lower global temperatures, snow and rain from that
evaporated water tends to be higher in oxygen-16, and the seawater
left behind tends to be higher in oxygen-18. Marine organisms then
incorporate more oxygen-18 into their skeletons and shells than they
would in a warmer climate. Paleoclimatologists also directly measure
this ratio in the water molecules of ice core samples as old as
hundreds of thousands of years.
Planetary geologists have measured the relative quantities of oxygen
isotopes in samples from the Earth, the Moon, Mars, and meteorites,
but were long unable to obtain reference values for the isotope ratios
in the Sun, believed to be the same as those of the primordial solar
nebula. Analysis of a silicon wafer exposed to the solar wind in space
and returned by the crashed Genesis spacecraft has shown that the Sun
has a higher proportion of oxygen-16 than does the Earth. The
measurement implies that an unknown process depleted oxygen-16 from
the Sun's disk of protoplanetary material prior to the coalescence of
dust grains that formed the Earth.
Oxygen presents two spectrophotometric absorption bands peaking at the
wavelengths 687 and 760 nm. Some remote sensing scientists have
proposed using the measurement of the radiance coming from vegetation
canopies in those bands to characterize plant health status from a
satellite platform. This approach exploits the fact that in those
bands it is possible to discriminate the vegetation's reflectance from
its fluorescence, which is much weaker. The measurement is technically
difficult owing to the low signal-to-noise ratio and the physical
structure of vegetation; but it has been proposed as a possible method
of monitoring the carbon cycle from satellites on a global scale.
Photosynthesis and respiration
================================
In nature, free oxygen is produced as a byproduct of light-driven
splitting of water during chlorophyllic photosynthesis. According to
some estimates, marine photoautotrophs such as red/green algae and
cyanobacteria provide about 70% of the free oxygen produced on Earth,
and the rest is produced in terrestrial environments by plants. 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.
A simplified overall formula for photosynthesis is
: 6 + 6 + photons → + 6
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. Many
steps are involved, but the result is the formation of a proton
gradient across the thylakoid membrane, which is used to synthesize
adenosine triphosphate (ATP) via photophosphorylation. The remaining
(after production of the water molecule) is released into the
atmosphere.
Oxygen is used in mitochondria of eukaryotes to generate ATP during
oxidative phosphorylation. The reaction for aerobic respiration is
essentially the reverse of photosynthesis and is simplified as
: + 6 → 6 + 6 + 2880 kJ/mol
In aquatic animals, dissolved oxygen in water is absorbed by gills,
through the skin or via the gut; in terrestrial animals such as
tetrapods, oxygen in air is actively taken into the body via lungs,
where gas exchange takes place to diffuse oxygen into the blood and
carbon dioxide out, and the body's circulatory system then transports
the oxygen to other tissues where cellular respiration takes place.
However in insects, the most successful and biodiverse terrestrial
clade, oxygen is directly conducted to the internal tissues via a deep
network of airways. Hemoglobin in red blood cells binds , changing
color from bluish red to bright red ( is released from another part of
hemoglobin through the Bohr effect). Other terrestrial invertebrates
use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders
and lobsters) instead. A liter of blood can dissolve up to 200 cm3 of
.
Until the discovery of anaerobic metazoa, oxygen was thought to be a
requirement for all complex life.
Reactive oxygen species, such as superoxide ion () and hydrogen
peroxide (), are reactive by-products of oxygen use in organisms.
Parts of the immune system of higher organisms 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. Oxygen is damaging to
obligately anaerobic organisms, which were the dominant form of early
life on Earth until began to accumulate in the atmosphere about 2.5
billion years ago during the Great Oxygenation Event, about a billion
years after the first appearance of these organisms.
An adult human at rest inhales 1.8 to 2.4 grams of oxygen per minute.
This amounts to more than 6 billion tonnes of oxygen inhaled by
humanity per year.
Living organisms
==================
Partial pressures of oxygen in the human body (PO2)
Unit !! Alveolar pulmonary gas pressures !! Arterial blood oxygen !!
Venous blood gas
kPa 14.2 11-13 4.0-5.3
mmHg 107 75-100 30-40
The free oxygen partial pressure in the body of a living vertebrate
organism is highest in the respiratory system, and decreases along any
arterial system, peripheral tissues, and venous system, respectively.
Partial pressure is the pressure that oxygen would have if it alone
occupied the volume.
Build-up in the atmosphere
============================
Free oxygen gas was almost nonexistent in Earth's atmosphere before
photosynthetic archaea and bacteria evolved, probably about 3.5
billion years ago. Free oxygen first appeared in significant
quantities during the Paleoproterozoic era (between 3.0 and 2.3
billion years ago). Even if there was much dissolved iron in the
oceans when oxygenic photosynthesis was getting more common, it
appears the banded iron formations were created by anoxyenic or
micro-aerophilic iron-oxidizing bacteria which dominated the deeper
areas of the photic zone, while oxygen-producing cyanobacteria covered
the shallows. Free oxygen began to outgas from the oceans 3-2.7
billion years ago, reaching 10% of its present level around 1.7
billion years ago.
The presence of large amounts of dissolved and free oxygen in the
oceans and atmosphere may have driven most of the extant anaerobic
organisms to extinction during the Great Oxygenation Event ('oxygen
catastrophe') about 2.4 billion years ago. Cellular respiration using
enables aerobic organisms to produce much more ATP than anaerobic
organisms. Cellular respiration of occurs in all eukaryotes,
including all complex multicellular organisms such as plants and
animals.
Since the beginning of the Cambrian period 540 million years ago,
atmospheric levels have fluctuated between 15% and 30% by volume.
Towards the end of the Carboniferous period (about 300 million years
ago) atmospheric levels reached a maximum of 35% by volume, which may
have contributed to the large size of insects and amphibians at this
time.
Variations in atmospheric oxygen concentration have shaped past
climates. When oxygen declined, atmospheric density dropped, which in
turn increased surface evaporation, causing precipitation increases
and warmer temperatures.
At the current rate of photosynthesis it would take about 2,000 years
to regenerate the entire in the present atmosphere.
It is estimated that oxygen on Earth will last for about one billion
years.
Extraterrestrial free oxygen
==============================
In the field of astrobiology and in the search for extraterrestrial
life oxygen is a strong biosignature. That said it might not be a
definite biosignature, being possibly produced abiotically on
celestial bodies with processes and conditions (such as a peculiar
hydrosphere) which allow free oxygen, like with Europa's and
Ganymede's thin oxygen atmospheres.
Industrial production
======================================================================
One hundred million tonnes of are extracted from air for industrial
uses annually by two primary methods. The most common method is
fractional distillation of liquefied air, with distilling as a vapor
while is left as a liquid.
The other primary method of producing is passing a stream of clean,
dry air through one bed of a pair of identical zeolite molecular
sieves, which absorbs the nitrogen and delivers a gas stream that is
90% to 93% . Simultaneously, nitrogen gas is released from the other
nitrogen-saturated zeolite bed, by reducing the chamber operating
pressure and diverting part of the oxygen gas from the producer bed
through it, in the reverse direction of flow. After a set cycle time
the operation of the two beds is interchanged, thereby allowing for a
continuous supply of gaseous oxygen to be pumped through a pipeline.
This is known as pressure swing adsorption. Oxygen gas is increasingly
obtained by these non-cryogenic technologies (see also the related
vacuum swing adsorption).
In academic laboratories, oxygen can be prepared by heating together
potassium chlorate mixed with a small proportion of manganese dioxide.
Oxygen gas can also be produced through electrolysis of water into
molecular oxygen and hydrogen. DC electricity must be used: if AC is
used, the gases in each limb consist of hydrogen and oxygen in the
explosive ratio 2:1. A similar method is the electrocatalytic
evolution from oxides and oxoacids. Chemical catalysts can be used as
well, such as in chemical oxygen generators or oxygen candles that are
used as part of the life-support equipment on submarines, and are
still part of standard equipment on commercial airliners in case of
depressurization emergencies. Another air separation method is forcing
air to dissolve through ceramic membranes based on zirconium dioxide
by either high pressure or an electric current, to produce nearly pure
gas.
Storage
======================================================================
Oxygen storage methods include high-pressure oxygen tanks, cryogenics
and chemical compounds. For reasons of economy, oxygen is often
transported in bulk as a liquid in specially insulated tankers, since
one liter of liquefied oxygen is equivalent to 840 liters of gaseous
oxygen at atmospheric pressure and 20 C. Such tankers are used to
refill bulk liquid-oxygen storage containers, which stand outside
hospitals and other institutions that need large volumes of pure
oxygen gas. Liquid oxygen is passed through heat exchangers, which
convert the cryogenic liquid into gas before it enters the building.
Oxygen is also stored and shipped in smaller cylinders containing the
compressed gas; a form that is useful in certain portable medical
applications and oxy-fuel welding and cutting.
Medical
=========
Uptake of 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 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, and 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.
Hyperbaric (high-pressure) medicine uses special oxygen chambers to
increase the partial pressure of around the patient and, when needed,
the medical staff. Carbon monoxide poisoning, gas gangrene, and
decompression sickness (the 'bends') are sometimes addressed with this
therapy. Increased 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 the blood. Increasing the pressure of as soon as possible helps to
redissolve the bubbles back into the blood so that these excess gasses
can be exhaled naturally through the lungs. Normobaric oxygen
administration at the highest available concentration is frequently
used as first aid for any diving injury that may involve inert gas
bubble formation in the tissues. There is epidemiological support for
its use from a statistical study of cases recorded in a long term
database.
Life support and recreational use
===================================
An application of as a low-pressure breathing gas is in modern space
suits, which surround their occupant's body with the breathing gas.
These devices use nearly pure oxygen at about one-third normal
pressure, resulting in a normal blood partial pressure of . This
trade-off of higher oxygen concentration for lower pressure is needed
to maintain suit flexibility.
Scuba and surface-supplied underwater divers and submarines also rely
on artificially delivered . Submarines, submersibles and atmospheric
diving suits usually operate at normal atmospheric pressure. Breathing
air is scrubbed of carbon dioxide by chemical extraction and oxygen is
replaced to maintain a constant partial pressure. Ambient pressure
divers breathe air or gas mixtures with an oxygen fraction suited to
the operating depth. Pure or nearly pure use in diving at pressures
higher than atmospheric is usually limited to rebreathers, or
decompression at relatively shallow depths (~6 meters depth, or less),
or medical treatment in recompression chambers at pressures up to 2.8
bar, where acute oxygen toxicity can be managed without the risk of
drowning. Deeper diving requires significant dilution of with other
gases, such as nitrogen or helium, to prevent oxygen toxicity.
People who climb mountains or fly in non-pressurized fixed-wing
aircraft sometimes have supplemental supplies. Pressurized commercial
airplanes have an emergency supply of automatically supplied to the
passengers 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
the sodium chlorate inside the canister. A steady stream of oxygen gas
is then produced by the exothermic reaction.
Oxygen, as a mild euphoric, has a history of recreational use in
oxygen bars and in sports. Oxygen bars are establishments found in the
United States since the late 1990s that offer higher than normal
exposure for a minimal fee. Professional athletes, especially in
American football, sometimes go off-field between plays to don oxygen
masks to boost performance. The pharmacological effect is doubted; a
placebo effect is a more likely explanation. Available studies support
a performance boost from oxygen enriched mixtures only if it is
inhaled 'during' aerobic exercise.
Other recreational uses that do not involve breathing include
pyrotechnic applications, such as George Goble's five-second ignition
of barbecue grills.
Industrial
============
Smelting of iron ore into steel consumes 55% of commercially produced
oxygen. In this process, is injected through a high-pressure lance
into molten iron, which removes sulfur impurities and excess carbon as
the respective oxides, and . The reactions are exothermic, so the
temperature increases to 1,700 °C.
Another 25% of commercially produced oxygen is used by the chemical
industry. Ethylene is reacted with to create ethylene oxide, which,
in turn, is converted into ethylene glycol; the primary feeder
material used to manufacture a host of products, including antifreeze
and polyester polymers (the precursors of many plastics and fabrics).
Most of the remaining 20% of commercially produced oxygen is used in
medical applications, metal cutting and welding, as an oxidizer in
rocket fuel, and in water treatment. Oxygen is used in oxyacetylene
welding, burning acetylene with to produce a very hot flame. In this
process, metal up to 60 cm thick is first heated with a small
oxy-acetylene flame and then quickly cut by a large stream of .
Compounds
======================================================================
The oxidation state of oxygen is −2 in almost all known compounds of
oxygen. The oxidation state −1 is found in a few compounds such as
peroxides. Compounds containing oxygen in other oxidation states are
very uncommon: −1/2 (superoxides), −1/3 (ozonides), 0 (elemental,
hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride), and
+2 (oxygen difluoride).
Oxides and other inorganic compounds
======================================
Water () is an oxide of hydrogen and the most familiar oxygen
compound. Hydrogen atoms are covalently bonded to oxygen in a water
molecule but also have an additional attraction (about 23.3 kJ/mol per
hydrogen atom) to an adjacent oxygen atom in a separate molecule.
These hydrogen bonds between water molecules hold them approximately
15% closer than what would be expected in a simple liquid with just
van der Waals forces.
Due to its electronegativity, oxygen forms chemical bonds with almost
all other elements to give corresponding oxides. The surface of most
metals, such as aluminium and titanium, are oxidized in the presence
of air and become coated with a thin film of oxide that passivates the
metal and slows further corrosion. Many oxides of the transition
metals are non-stoichiometric compounds, with slightly less metal than
the chemical formula would show. For example, the mineral FeO
(wüstite) is written as , where 'x' is usually around 0.05.
Oxygen is present in the atmosphere in trace quantities in the form of
carbon dioxide (). The Earth's crustal rock is composed in large part
of oxides of silicon (silica , as found in granite and quartz),
aluminium (aluminium oxide , in bauxite and corundum), iron (iron(III)
oxide , in hematite and rust), and calcium carbonate (in limestone).
The rest of the Earth's crust is also made of oxygen compounds, in
particular various complex silicates (in silicate minerals). The
Earth's mantle, of much larger mass than the crust, is largely
composed of silicates of magnesium and iron.
Water-soluble silicates in the form of , , and are used as detergents
and adhesives.
Oxygen also acts as a ligand for transition metals, forming transition
metal dioxygen complexes, which feature metal-. This class of
compounds includes the heme proteins hemoglobin and myoglobin. An
exotic and unusual reaction occurs with platinum hexafluoride, which
oxidizes oxygen to give , dioxygenyl hexafluoroplatinate.
Organic compounds
===================
Among the most important classes of organic compounds that contain
oxygen are (where "R" is an organic group): alcohols (); ethers ();
ketones (); aldehydes (); carboxylic acids (); esters (); acid
anhydrides (); and amides (). There are many important organic
solvents that contain oxygen, including: acetone, methanol, ethanol,
isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF,
DMSO, acetic acid, and formic acid. Acetone () and phenol () are used
as feeder materials in the synthesis of many different substances.
Other important organic compounds that contain oxygen are: glycerol,
formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and
acetamide. Epoxides are ethers in which the oxygen atom is part of a
ring of three atoms. The element is similarly found in almost all
biomolecules that are important to (or generated by) life.
Oxygen reacts spontaneously with many organic compounds at or below
room temperature in a process called autoxidation. Most of the organic
compounds that contain oxygen are not made by direct action of .
Organic compounds important in industry and commerce that are made by
direct oxidation of a precursor include ethylene oxide and peracetic
acid.
Safety and precautions
======================================================================
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The NFPA 704 standard rates compressed oxygen gas as nonhazardous to
health, nonflammable and nonreactive, but an oxidizer. Refrigerated
liquid oxygen (LOX) is given a health hazard rating of 3 (for
increased risk of hyperoxia from condensed vapors, and for hazards
common to cryogenic liquids such as frostbite), and all other ratings
are the same as the compressed gas form.
Toxicity
==========
Oxygen gas () can be toxic at elevated partial pressures, leading to
convulsions and other health problems. Oxygen toxicity usually begins
to occur at partial pressures more than 50 kilopascals (kPa), equal to
about 50% oxygen composition at standard pressure or 2.5 times the
normal sea-level partial pressure of about 21 kPa. This is not a
problem except for patients on mechanical ventilators, since gas
supplied through oxygen masks in medical applications is typically
composed of only 30-50% by volume (about 30 kPa at standard
pressure).
At one time, premature babies were placed in incubators containing
-rich air, but this practice was discontinued after some babies were
blinded by the oxygen content being too high.
Breathing pure in space applications, such as in some modern space
suits, or in early spacecraft such as Apollo, causes no damage due to
the low total pressures used. In the case of spacesuits, the partial
pressure in the breathing gas is, in general, about 30 kPa (1.4 times
normal), and the resulting partial pressure in the astronaut's
arterial blood is only marginally more than normal sea-level partial
pressure.
Oxygen toxicity to the lungs and central nervous system can also occur
in deep scuba diving and surface-supplied diving. Prolonged breathing
of an air mixture with an partial pressure more than 60 kPa can
eventually lead to permanent pulmonary fibrosis. Exposure to an
partial pressure greater than 160 kPa (about 1.6 atm) may lead to
convulsions (normally fatal for divers). Acute oxygen toxicity
(causing seizures, its most feared effect for divers) can occur by
breathing an air mixture with 21% at 66 m or more of depth; the same
thing can occur by breathing 100% at only 6 m.
Combustion and other hazards
==============================
Highly concentrated sources of oxygen promote rapid combustion. Fire
and explosion hazards exist when concentrated oxidants and fuels are
brought into close proximity; an ignition event, such as heat or a
spark, is needed to trigger combustion. Oxygen is the oxidant, not the
fuel.
Concentrated will allow combustion to proceed rapidly and
energetically. Steel pipes and storage vessels used to store and
transmit both gaseous and liquid oxygen will act as a fuel; and
therefore the design and manufacture of systems requires special
training to ensure that ignition sources are minimized. The fire that
killed the Apollo 1 crew in a launch pad test spread so rapidly
because the capsule was pressurized with pure but at slightly more
than atmospheric pressure, instead of the normal pressure that would
be used in a mission.
Liquid oxygen spills, if allowed to soak into organic matter, such as
wood, petrochemicals, and asphalt can cause these materials to
detonate unpredictably on subsequent mechanical impact.
See also
======================================================================
* Geological history of oxygen
* Hypoxia (environmental) for depletion in aquatic ecology
* Ocean deoxygenation
* Hypoxia (medical), a lack of oxygen
* Limiting oxygen concentration
* Oxygen compounds
* Oxygen plant
* Oxygen sensor
* Dark oxygen
External links
======================================================================
* [
https://www.periodicvideos.com/videos/008.htm Oxygen] at 'The
Periodic Table of Videos' (University of Nottingham)
* [
https://www.organic-chemistry.org/chemicals/oxidations/oxygen.shtm
Oxidizing Agents > Oxygen]
* [
https://www.uigi.com/oxygen.html Oxygen (O2) Properties, Uses,
Applications]
* [
https://www.americanscientist.org/issues/pub/the-story-of-o Roald
Hoffmann article on "The Story of O"]
* [
https://www.webelements.com/webelements/elements/text/O/index.html
WebElements.com - Oxygen]
*
* [
https://scrippso2.ucsd.edu/ Scripps Institute: Atmospheric Oxygen
has been dropping for 20 years]
License
=========
All content on Gopherpedia comes from Wikipedia, and is licensed under CC-BY-SA
License URL:
http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Oxygen
