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Essay/Term paper: Oxygen

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Oxygen


Oxygen and its compounds play a key role in many of the important processes of
life and industry. Oxygen in the biosphere is essential in the processes of
respiration and metabolism, the means by which animals derive the energy needed
to sustain life. Furthermore, oxygen is the most abundant element at the surface
of the Earth. In combined form it is found in ores, earths, rocks, and gemstones,
as well as in all living organisms. Oxygen is a gaseous chemical element in
Group VA of the periodic table. The chemical symbol for atomic oxygen is O, its
atomic number is 8, and its atomic weight is 15.9994. Elemental oxygen is known
principally in the gaseous form as the diatomic molecule, which makes up 20.95%
of the volume of dry air. Diatomic oxygen is colorless, odorless, and tasteless.
Two 18th-century scientists share the credit for first isolating elemental
oxygen: Joseph PRIESTLEY (1733-1804), an English clergyman who was employed as a
literary companion to Lord Shelburne at the time of his most significant
experimental work, and Carl Wilhelm SCHEELE (1742-86), a Swedish pharmacist and
chemist. It is generally believed that Scheele was the first to isolate oxygen,
but that Priestley, who independently achieved the isolation of oxygen somewhat
later, was the first to publicly announce his findings. The interpretation of
the findings of Priestley and the resultant clarification of the nature of
oxygen as an element was accomplished by the French scientist Antoine-Laurent
LAVOISIER (1743-94). Lavoisier's experimental work, which extended and improved
upon Priestley's experiments, was principally responsible for the understanding
of COMBUSTION and the establishment of the law of conservation of matter.
Lavoisier gave oxygen its name, which is derived from two Greek words that mean
"acid former." Lavoisier held the mistaken idea that oxides, when dissolved in
water, would form only acids. It is true that some oxides when dissolved in
water do form acids; for example, sulfur dioxide forms sulfurous acid. Some
oxides, however, such as sodium oxide, dissolve in water to form bases, as in
the reaction to form sodium hydroxide; therefore oxygen was actually
inappropriately named.

NATURAL OCCURRENCE

Oxygen is formed by a number of nuclear processes that are believed to occur in
stellar interiors. The most abundant isotope of oxygen, with mass 16, is thought
to be formed in hydrogen-burning stars by the capture of a proton by the
isotopes of nitrogen and fluorine, with the subsequent emission of, respectively,
a gamma ray and an alpha particle. In helium-burning stars the isotope of carbon
with mass 12 is thought to capture an alpha particle to form the isotope with
mass 16 with the emission of a gamma ray. In the terrestrial environment oxygen
accounts for about half of the mass of the Earth's crust, 89% of the mass of the
oceans, and 23% of the mass (and 21% of the volume) of the atmosphere. Most of
the Earth's rocks and soils are principally silicates. The silicates are an
amazingly complex group of materials that typically consist of greater than 50
(atomic) percent oxygen in combination with silicon and one or more metallic
elements. Several important ores are principally oxides of the desired metals,
such as the important iron-bearing minerals hematite, magnetite, and limonite
and the most important aluminum-bearing mineral, BAUXITE (a mixture of hydrated
aluminum oxides and iron oxide).

PHYSICAL AND CHEMICAL PROPERTIES

Three naturally occurring isotopes of oxygen have been found: one with mass 16
(99. 759% of all oxygen), one with mass 17 (0.037%); and one with mass 18
(0.204%). The rarer isotopes, principally the latter, find their major use in
labeling experiments used by scientists to follow the steps of chemical
reactions. If oxygen at a pressure of one atmosphere is cooled, it will liquefy
at 90.18 K (-182.97 deg C; -297.35 deg F), the normal boiling point of oxygen,
and it will solidify at 54.39 K (-218.76 deg C; -361.77 deg F), the normal
melting point of oxygen. The liquid and solid forms of oxygen have a pale blue
color. Several different structures are known for solid oxygen: solid type III,
from the lowest temperatures achievable to 23.66 K; type II, from 23.66 to 43.76
K; and type I, from 43.76 to 54.39 K. The critical temperature for oxygen, the
temperature above which it is impossible to liquefy the gas no matter how much
pressure is applied, is 154.3 K (-118.9 deg C; -181.9 deg F). The pressure of
liquid and gaseous oxygen coexisting in equilibrium at the critical temperature
is 49.7 atmospheres. Oxygen gas exhibits a slight but important solubility in
water. Molecular oxygen dissolved in water is required by aquatic organisms for
their metabolic processes and is ultimately responsible for the oxidation and
removal of organic wastes in water. The solubilities of gases depend on the
temperature of the solution and the pressure of the gas over the solution. At 20
deg C (68 deg F) and an oxygen pressure of one atmosphere, the solubility of
O(2) in water is about 45 grams of oxygen per cubic meter of water, or 45 ppm
(parts per million). Molecular diatomic oxygen is a fairly stable molecule
requiring a dissociation energy (the energy required to dissociate one mole of
molecular oxygen in its ground state into two moles of atomic oxygen in its
ground state) of 493.6 kilojoules per mole. The molecule is dissociated by
ultraviolet radiation of any wavelength shorter than 193 nm. Solar radiation
striking stratospheric oxygen dissociates it into atomic oxygen for this reason.
The atomic oxygen formed in this fashion is capable of reacting with oxygen to
form OZONE.

Corrosion

Many direct, uncatalyzed reactions of oxygen do not occur rapidly at room
temperature. This fact has a number of important consequences. One of these
consequences has to do with the use of metals as structural materials. Metals
that are used in construction, such as iron (principally as steel) and aluminum,
form highly stable oxides. For example, the oxidation of aluminum has a
significant tendency to occur. However, in spite of this tendency, the reaction
occurs so slowly at room temperature that it can be said for most practical
purposes not to occur at all, and for this reason aluminum is an appropriate and
widely used structural material. The slowness of this reaction is due in part to
the stability of the oxygen-oxygen bond and in part because of a very thin,
protective layer of oxide that forms on the surface of the aluminum. The
oxidation of iron is a complex process involving impurities in the iron, as well
as water and carbon dioxide. This oxidative destruction, or rusting, of iron and
steel--which are among our most important structural materials--is extremely
costly to modern societies.

Biological Oxidation

Another important aspect of the rates of oxygen reactions concerns the rate of
reaction with organic materials. Such oxidation reactions are, ultimately, the
sources of energy for the higher plants and animals, are responsible for the
cleansing of streams of biodegradable wastes, and are responsible for the
natural decomposition of organic material. The rates of reactions in this
category are selectively controlled by enzymes in the organisms that facilitate
the reactions. Thus waste products and dead plants and animals decompose (are
oxidized) principally through the agency of microorganisms, and energy-bearing
foods are metabolized (oxidized) by means of biological processes.

Reactivity

There is a marked difference between the rates of reactions with oxygen at room
temperature and the rates at elevated temperatures. Many substances that do not
react rapidly with oxygen in air at temperatures below 100 deg C will do so at
1000 deg C with a strong evolution of heat (exothermically). For example, coal
and petroleum can be stored indefinitely at the temperatures encountered under
normal climatic conditions, but they readily oxidize, exothermically, at
elevated temperatures. The most common compounds of oxygen are those in which
the element exhibits a valence of two. This fact is associated with the
electronic structure of atomic oxygen; this atom requires two additional
electrons to fill its outermost energy level. Examples of divalent oxides are
numerous among well-known substances such as water; carbon dioxide; aluminum
oxide; silicon dioxide; the silicates, calcium carbonate or limestone; and
sulfur dioxide. Oxygen is also known to have other valences, such as in the
PEROXIDES, of which hydrogen peroxide is an example. The direct reaction of
oxygen with another element frequently follows the pattern discussed above; that
is, it does not occur rapidly or at all at room temperature but is strongly
exothermic, and once oxidation is initiated the evolved heat raises the
temperature of the reactants such that the reaction is self-sustaining. Examples
of such reactions are with the elements magnesium, carbon, and hydrogen.
Magnesium and carbon burn in air once the reaction is initiated, and a hydrogen-
oxygen mixture can react explosively when the reaction is initiated by a flame
or spark. The explosion of a hydrogen-oxygen mixture is an extremely fast
reaction and occurs because of the formation of atomic oxygen in the exploding
mixture.

Uses

Pure oxygen is used extensively in technological processes. It is used in the
welding, cutting, and forming of metals, as in oxyacetylene welding, in which
oxygen reacts with acetylene to form an extremely hot flame. Oxygen is added to
the inlet air (3 to 5%) in modern blast furnaces to increase the temperature in
the furnace; it is also used in the basic oxygen converter for steel production,
in the manufacture of chemicals, and for rocket propulsion. Oxygen is also used
in the partial combustion of methane (natural gas) or coal (taken here to be
carbon) to make mixtures of carbon monoxide and hydrogen called synthesis gas,
which is in turn used for the manufacture of methanol. Processes in which
combustible liquids are produced from coal will become increasingly important as
petroleum resources become further depleted.

PRODUCTION

Oxygen is conveniently produced in the laboratory by heating mercuric oxide or
potassium chlorate to moderately high temperatures. The production from mercuric
oxide is the method that was employed by Joseph Priestley, and the production
from the potassium chlorate method commonly used by students in today's
laboratories. Oxygen is liberated when solid potassium chlorate is heated to 400
deg C or, when manganese dioxide is added as a CATALYST, to 200 deg C. The
liberated oxygen can be collected by water displacement because of the low
solubility of oxygen in water. Oxygen can also be produced in the laboratory by
the electrolysis of water, a process that reverses the violent hydrogen-oxygen
reaction discussed previously. When a current is passed through water the liquid
is decomposed at the electrodes. This method is also used to produce oxygen on a
commercial scale when a high-purity product is desired. The more economical, and
therefore preferred, method for the commercial production of oxygen is the
liquefaction and distillation of air. The air is cooled until it liquefies,
principally by being made to do work in a rotating expansion turbine, and the
resulting liquid air is fractionated by a complex distillation process. The
gaseous oxygen produced in this fashion is shipped in pressurized cylinders or,
as is often the case when large amounts are involved, through pipelines to
nearby industrial plants.

RELATIONSHIP TO LIFE SCIENCES

Most organisms depend on oxygen to sustain their biological processes. The great
majority of living organisms fall into two categories. In the first category are
the higher plants and the photosynthetic bacteria. These organisms utilize light
energy through PHOTOSYNTHESIS to combine carbon dioxide and water (or,
infrequently, other inorganic substances in place of water) into more complex
materials characterized as CARBOHYDRATES while at the same time releasing oxygen
into the atmosphere. In the second category are the higher animals, most
microorganisms, and photosynthetic cells that live in the dark. All these
second-category organisms use complex series of enzyme-catalyzed OXIDATION AND
REDUCTION reactions using materials such as glucose as the fuel and oxygen as
the terminal oxidizing agent (see METABOLISM). The end products of metabolism in
these organisms are carbon dioxide and water, which are returned to the
atmosphere. The net result of these complementary functions is the oxygen cycle,
in which the photosynthetic organisms, using solar energy, synthesize
carbohydrates from water and carbon dioxide and give off oxygen as a by-product,
while the aerobic organisms oxidize ingested organic materials, using up oxygen
and giving off carbon dioxide and water through a complex series of metabolic
processes. It has been estimated that 3.5 X (10 to the power of 11) tons of
carbon dioxide are cycled annually via these processes. Thus, in the
vertebrates--and in humans in particular--oxygen is necessary to sustain
metabolism and thus life. Air is inhaled and oxygen in the air is exchanged in
the lungs between the atmosphere and the hemoglobin in the blood. The blood
carries the oxygen, complexed with hemoglobin, to all parts of the body in which
metabolic processes occur. It also carries carbon dioxide back to the lungs,
where the carbon dioxide is exchanged with the atmosphere and exhaled. If the
oxygen concentration were to drop to about half its value in the atmosphere,
humans could no longer survive. For this reason an important component of the
life-support systems of divers and astronauts is a source of oxygen gas.
Similarly, persons ill with respiratory diseases that interfere with normal
respiration, such as pneumonia and emphysema, are often kept in OXYGEN TENTS and
HYPERBARIC CHAMBERS, the latter administering high-pressure oxygen, may be used
to treat a variety of ailments.

IMPORTANT COMPOUNDS

In the realm of inorganic chemistry there is a very large number of oxygen-
containing compounds. There are very few elements for which no oxides are known,
and there are several metallic elements (such as titanium, vanadium, and
praseodymium) for which a wide variety of solid oxides exist. The solid oxides
of the metallic elements can generally be synthesized by the direct reaction of
the elements at high temperatures. In many cases such reactions will result in
the formation of a single oxide of the metal in its most oxidized form. Typical
examples are the metallic oxides of sodium, calcium, lanthanum, titanium,
vanadium, and tungsten. In the cases of elements capable of forming reduced
oxides, in particular the early transition metals, the reduced oxides can be
formed by heating the highest oxide, formed as above, to very high temperatures
(1,500 K or higher) either in an inert container or in the presence of the
metallic element. The reduced oxides that result exhibit a variation in the
extent and importance of direct metal-metal bonding in the compounds, and this
variation gives rise to a variety of electrical and magnetic properties. The
more metal-rich of these oxides are metallic conductors and tend to be
nonstoichiometric; that is, they are observed to exist over a range of
compositions all possessing the same underlying structure. A number of these
titanium oxides exhibit more than one crystal structure (polymorphism). The most
oxidized compound, titanium, is widely used in the RUTILE form as a white
pigment in paints. Ternary oxides, consisting of two metallic elements plus
oxygen, are of great interest to solid-state scientists. For example, compounds
such as the SPINELS and the PEROVSKITES are studied extensively because of their
interesting magnetic and electrical properties. Examples of important ternary
oxides are the magnetic FERRITES, whose magnetic properties can be tailored,
making them useful in computer memory units. The ferrites are prepared by firing
compacted mixtures of iron oxide and one or more metal oxides (such as those of
nickel, copper, zinc, magnesium, and manganese). Also of importance in inorganic
chemistry are the oxides of the nonmetals. Most of the nonmetals are known to
form a wide variety of compounds with oxygen. The nitrogen oxides are
undesirable by-products of high-temperature combustion in air (as in an internal
combustion engine) and can cause serious environmental pollution.


 

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