water | Definition, Chemical Formula, Structure and more

A number of natural states of water exist. It forms precipitation in the form of rain and aerosols in the form of fog. Clouds consist of suspended droplets of water and ice, its solid state. When finely divided, crystalline ice may precipitate in the form of snow. The gaseous state of water is steam or water vapor.


Define Water in detail

water, a substance composed of the chemical elements hydrogen and oxygen and existing in gaseous, liquid, and solid states. It is one of the most plentiful and essential of compounds. A tasteless and odourless liquid at room temperature, it has the important ability to dissolve many other substances. Indeed, the versatility of water as a solvent is essential to living organisms. Life is believed to have originated in the aqueous solutions of the world’s oceans, and living organisms depend on aqueous solutions, such as blood and digestive juices, for biological processes. Water also exists on other planets and moons both within and beyond the solar system. In small quantities water appears colourless, but water actually has an intrinsic blue colour caused by slight absorption of light at red wavelengths.

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Although the molecules of water are simple in structure (H2O), the physical and chemical properties of the compound are extraordinarily complicated, and they are not typical of most substances found on Earth. For example, although the sight of ice cubes floating in a glass of ice water is commonplace, such behaviour is unusual for chemical entities. For almost every other compound, the solid state is denser than the liquid state; thus, the solid would sink to the bottom of the liquid. The fact that ice floats on water is exceedingly important in the natural world, because the ice that forms on ponds and lakes in cold areas of the world acts as an insulating barrier that protects the aquatic life below. If ice were denser than liquid water, ice forming on a pond would sink, thereby exposing more water to the cold temperature. Thus, the pond would eventually freeze throughout, killing all the life-forms present.

Ice cubes on white background. (frozen; freeze; ice cube)
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Water occurs as a liquid on the surface of Earth under normal conditions, which makes it invaluable for transportation, for recreation, and as a habitat for a myriad of plants and animals. The fact that water is readily changed to a vapour (gas) allows it to be transported through the atmosphere from the oceans to inland areas where it condenses and, as rain, nourishes plant and animal life. (See hydrosphere: The hydrologic cycle for a description of the cycle by which water is transferred over Earth.)

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Because of its prominence, water has long played an important religious and philosophical role in human history. In the 6th century BCE, Thales of Miletus, sometimes credited for initiating Greek philosophy, regarded water as the sole fundamental building block of matter:

It is water that, in taking different forms, constitutes the earth, atmosphere, sky, mountains, gods and men, beasts and birds, grass and trees, and animals down to worms, flies and ants. All these are different forms of water. Meditate on water!

Two hundred years later, Aristotle considered water to be one of four fundamental elements, in addition to earth, air, and fire. The belief that water was a fundamental substance persisted for more than 2,000 years until experiments in the second half of the 18th century showed that water is a compound made up of the elements hydrogen and oxygen.

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The water on the surface of Earth is found mainly in its oceans (97.25 percent) and polar ice caps and glaciers (2.05 percent), with the balance in freshwater lakes, rivers, and groundwater. As Earth’s population grows and the demand for fresh water increases, water purification and recycling become increasingly important. Interestingly, the purity requirements of water for industrial use often exceed those for human consumption. For example, the water used in high-pressure boilers must be at least 99.999998 percent pure. Because seawater contains large quantities of dissolved salts, it must be desalinated for most uses, including human consumption.

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This article describes the molecular structure of water as well as its physical and chemical properties. For other major treatments of water, see climate; environmental works; hydrosphere; ice; and pollution.

Structure of water

Liquid water

The water molecule is composed of two hydrogen atoms, each linked by a single chemical bond to an oxygen atom. Most hydrogen atoms have a nucleus consisting solely of a proton. Two isotopic forms, deuterium and tritium, in which the atomic nuclei also contain one and two neutrons, respectively, are found to a small degree in water. Deuterium oxide (D2O), called heavy water, is important in chemical research and is also used as a neutron moderator in some nuclear reactors.

Although its formula (H2O) seems simple, water exhibits very complex chemical and physical properties. For example, its melting point, 0 °C (32 °F), and boiling point, 100 °C (212 °F), are much higher than would be expected by comparison with analogous compounds, such as hydrogen sulfide and ammonia. In its solid form, ice, water is less dense than when it is liquid, another unusual property. The root of these anomalies lies in the electronic structure of the water molecule.

The water molecule is not linear but bent in a special way. The two hydrogen atoms are bound to the oxygen atom at an angle of 104.5°.structure of the water molecule showing the two hydrogen atoms bound to the oxygen atom at an angle of 104.5 degrees.

The O―H distance (bond length) is 95.7 picometres (9.57 × 10−11 metres, or 3.77 × 10−9 inches). Because an oxygen atom has a greater electronegativity than a hydrogen atom, the O―H bonds in the water molecule are polar, with the oxygen bearing a partial negative charge (δ−) and the hydrogens having a partial positive charge (δ+).structure of the water molecule showing the charges of the hydrogen and oxygen atoms

Hydrogen atoms in water molecules are attracted to regions of high electron density and can form weak linkages, called hydrogen bonds, with those regions. This means that the hydrogen atoms in one water molecule are attracted to the nonbonding electron pairs of the oxygen atom on an adjacent water molecule. The structure of liquid water is believed to consist of aggregates of water molecules that form and re-form continually. This short-range order, as it is called, accounts for other unusual properties of water, such as its high viscosity and surface tension.

An oxygen atom has six electrons in its outer (valence) shell, which can hold a total of eight electrons. When an oxygen atom forms a single chemical bond, it shares one of its own electrons with the nucleus of another atom and receives in return a share of an electron from that atom. When bonded to two hydrogen atoms, the outer electron shell of the oxygen atom is filled.

The electron arrangement in the water molecule can be represented as follows.structure of the water molecule showing the electron arrangement

Each pair of dots represents a pair of unshared electrons (i.e., the electrons reside on only the oxygen atom). This situation can also be depicted by placing the water molecule in a cube.

Water molecule in a cube showing unshared electrons. Hydrogen bonding.

Each ↑↓ symbol represents a pair of unshared electrons. This electronic structure leads to hydrogen bonding.structure of a water molecule showing unshared electrons, leading to hydrogen bonding

Structures of ice by water

In the solid state (ice), intermolecular interactions lead to a highly ordered but loose structure in which each oxygen atom is surrounded by four hydrogen atoms; two of these hydrogen atoms are covalently bonded to the oxygen atom, and the two others (at longer distances) are hydrogen bonded to the oxygen atom’s unshared electron pairs.

This open structure of ice causes its density to be less than that of the liquid state, in which the ordered structure is partially broken down and the water molecules are (on average) closer together. When water freezes, a variety of structures are possible depending on the conditions. Eighteen different forms of ice are known and can be interchanged by varying external pressure and temperature.

Significance of the structure of liquid water

The liquid state of water has a very complex structure, which undoubtedly involves considerable association of the molecules. The extensive hydrogen bonding among the molecules in liquid water produces much larger values for properties such as viscosity, surface tension, and boiling point than are expected for a typical liquid containing small molecules. For example, based on the size of its molecules, water would be expected to have a boiling point nearly 200 °C (360 °F) lower than its observed boiling point. In contrast to the condensed states (solid and liquid) of water, which exhibit extensive association among the water molecules, its gaseous (vapour) phase contains relatively independent water molecules at large distances from each other.

The polarity of the water molecule plays a major part in the dissolution of ionic compounds during the formation of aqueous solutions. Earth’s oceans contain vast amounts of dissolved salts, which provide a great natural resource. In addition, the hundreds of chemical reactions that occur every instant to keep organisms alive all take place in aqueous fluids. Also, the ability of foods to be flavoured as they are cooked is made possible by the solubility in water of such substances as sugar and salt. Although the solubility of substances in water is an extremely complex process, the interaction between the polar water molecules and the solute (i.e., the substance being dissolved) plays a major role. When an ionic solid dissolves in water, the positive ends of the water molecules are attracted to the anions, while their negative ends are attracted to the cations. This process is called hydration. The hydration of its ions tends to cause a salt to break apart (dissolve) in the water. In the dissolving process the strong forces present between the positive and negative ions of the solid are replaced by strong water-ion interactions.

When ionic substances dissolve in water, they break apart into individual cations and anions. For instance, when sodium chloride (NaCl) dissolves in water, the resulting solution contains separated Na+ and Cl ions.formula representing when sodium chloride (NaCl) dissolves in water, the resulting solution contains separated positive Na and Cl ions.

In this equation the (s) represents the solid state, and the (aq), which is an abbreviation for aqueous, shows that the ions are hydrated—that is, they have a certain number of water molecules attached to them. As sodium chloride dissolves, four water molecules closely associate with the sodium ion. (The hydration number of Na+ is four.) Just outside this inner hydration sphere is a region where water molecules are partially ordered by the presence of the [Na(H2O)4]+ hydrated ion. This partially ordered region blends into “regular” (bulk) liquid water.

Generally speaking, the greater the charge density (the ratio of charge to surface area) of an ion, the larger the hydration number will be. As a rule, negative ions have smaller hydration numbers than positive ions because of the greater crowding that occurs when the hydrogen atoms of the water molecules are oriented toward the anion.

Many nonionic compounds are also soluble in water. For example, ethanol (C2H5OH), the alcoholic component of wine, beer, and distilled spirits, is highly soluble in water. These beverages contain varying percentages of ethanol in aqueous solution with other substances. Ethanol is so soluble in water because of the structure of the alcohol molecule. The molecule contains a polar O―H bond like those in water, which allows it to interact effectively with water.

There are many substances for which water is not an acceptable solvent. Animal fat, for example, is insoluble in pure water because the nonpolar nature of fat molecules renders them incompatible with polar water molecules. In general, polar and ionic substances are soluble in water. A useful rule of thumb for determining whether two substances are likely to be miscible (i.e., will mix to form a solution) is “like dissolves like.” That is, two polar substances are likely to mix to form a solution, as are two nonpolar substances.

Behaviour and properties

Water at high temperatures and pressures

The characteristic ability of water to behave as a polar solvent (dissolving medium) changes when water is subjected to high temperatures and pressures. As water becomes hotter, the molecules seem much more likely to interact with nonpolar molecules. For example, at 300 °C (572 °F) and high pressure, water has dissolving properties very similar to acetone (CH3COCH3), a common organic solvent.

Water exhibits particularly unusual behaviour beyond its critical temperature and pressure (374 °C [705.2 °F], 218 atmospheres). Above its critical temperature, the distinction between the liquid and gaseous states of water disappears—it becomes a supercritical fluid, the density of which can be varied from liquidlike to gaslike by varying its temperature and pressure. If the density of supercritical water is high enough, ionic solutes are readily soluble, as is true for “normal” water; but, surprisingly, this supercritical fluid can also readily dissolve nonpolar substances—something ordinary water cannot do. Because of its ability to dissolve nonpolar substances, supercritical water can be used as a combustion medium for destroying toxic wastes. For example, organic wastes can be mixed with oxygen in sufficiently dense supercritical water and combusted in the fluid; the flame actually burns “underwater.” Oxidation in supercritical water can be used to destroy a wide variety of hazardous organic substances with the advantage that a supercritical-water reactor is a closed system, so there are no emissions released into the atmosphere.

Physical properties of water

Water has several important physical properties. Although these properties are familiar because of the omnipresence of water, most of the physical properties of water are quite atypical. Given the low molar mass of its constituent molecules, water has unusually large values of viscosity, surface tension, heat of vaporization, and entropy of vaporization, all of which can be ascribed to the extensive hydrogen bonding interactions present in liquid water. The open structure of ice that allows for maximum hydrogen bonding explains why solid water is less dense than liquid water—a highly unusual situation among common substances.

Selected physical properties of water
molar mass 18.0151 grams per mole
melting point 0.00 °C
boiling point 100.00 °C
maximum density (at 3.98 °C) 1.0000 grams per cubic centimetre
density (25 °C) 0.99701 grams per cubic centimetre
vapour pressure (25 °C) 23.75 torr
heat of fusion (0 °C) 6.010 kilojoules per mole
heat of vaporization (100 °C) 40.65 kilojoules per mole
heat of formation (25 °C) −285.85 kilojoules per mole
entropy of vaporization (25 °C) 118.8 joules per °C mole
viscosity 0.8903 centipoise
surface tension (25 °C) 71.97 dynes per centimeter

Chemical properties of water

Acid-base reactions

Water undergoes various types of chemical reactions. One of the most important chemical properties of water is its ability to behave as both an acid (a proton donor) and a base (a proton acceptor), the characteristic property of amphoteric substances. This behaviour is most clearly seen in the autoionization of water:H2O(l) + H2O(l) ⇌ H3O+(aq) + OH(aq),where the (l) represents the liquid state, the (aq) indicates that the species are dissolved in water, and the double arrows indicate that the reaction can occur in either direction and an equilibrium condition exists. At 25 °C (77 °F) the concentration of hydrated H+ (i.e., H3O+, known as the hydronium ion) in water is 1.0 × 10−7 M, where M represents moles per litre. Since one OH ion is produced for each H3O+ ion, the concentration of OH at 25 °C is also 1.0 × 10−7 M. In water at 25 °C the H3O+ concentration and the OH concentration must always be 1.0 × 10−14:[H+][OH] = 1.0 × 10−14,where [H+] represents the concentration of hydrated H+ ions in moles per litre and [OH] represents the concentration of OH ions in moles per litre.

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When an acid (a substance that can produce H+ ions) is dissolved in water, both the acid and the water contribute H+ ions to the solution. This leads to a situation in which the H+ concentration is greater than 1.0 × 10−7 M. Since it must always be true that [H+][OH] = 1.0 × 10−14 at 25 °C, the [OH] must be lowered to some value below 1.0 × 10−7. The mechanism for reducing the concentration of OH involves the reactionH+ + OH → H2O,which occurs to the extent needed to restore the product of [H+] and [OH] to 1.0 × 10−14 M. Thus, when an acid is added to water, the resulting solution contains more H+ than OH; that is, [H+] > [OH]. Such a solution (in which [H+] > [OH]) is said to be acidic.

The most common method for specifying the acidity of a solution is its pH, which is defined in terms of the hydrogen ion concentration:pH = −log [H+],where the symbol log stands for a base-10 logarithm. In pure water, in which [H+] = 1.0 × 10−7 M, the pH = 7.0. For an acidic solution, the pH is less than 7. When a base (a substance that behaves as a proton acceptor) is dissolved in water, the H+ concentration is decreased so that [OH] > [H+]. A basic solution is characterized by having a pH > 7. In summary, in aqueous solutions at 25 °C:

neutral solution [H+] = [OH] pH = 7
acidic solution [H+] > [OH] pH < 7
basic solution [OH] > [H+] pH > 7

Oxidation-reduction reactions

When an active metal such as sodium is placed in contact with liquid water, a violent exothermic (heat-producing) reaction occurs that releases flaming hydrogen gas.2Na(s) + 2H2O(l) → 2Na+(aq) + 2OH(aq) + H2(g)This is an example of an oxidation-reduction reaction, which is a reaction in which electrons are transferred from one atom to another. In this case, electrons are transferred from sodium atoms (forming Na+ ions) to water molecules to produce hydrogen gas and OH ions. The other alkali metals give similar reactions with water. Less-active metals react slowly with water. For example, iron reacts at a negligible rate with liquid water but reacts much more rapidly with superheated steam to form iron oxide and hydrogen gas.Water: formula describing iron reacts at a negligible rate with liquid water but reacts much more rapidly with superheated steam to form iron oxide and hydrogen gas.

Noble metals, such as gold and silver, do not react with water at all.


zirconia, zirconium dioxide, an industrially important compound of zirconium and oxygen usually derived from the mineral zircon


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waterpower, power produced by a stream of water as it turns a wheel or similar device. The waterwheel was probably invented in the 1st century BCE, and it was widely used throughout the Middle Ages and into modern times for grinding grain, operating bellows for furnaces, and other purposes. The more-compact water turbine, which passes water through a series of fixed and rotating blades, was introduced in 1827 by Benoît Fourneyron, a French experimenter, whose first turbine developed about 6 horsepower. By 1832 he had perfected a turbine capable of developing 50 horsepower. Various modifications followed Fourneyron’s design, notably those of James Thomson (about 1851) and James B. Francis (1855), using radial flow inward. Water turbines, used originally for direct mechanical drive for irrigation, now are used almost exclusively to generate electric power. See also hydroelectric power.


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anomalous water


anomalous water, also called Orthowater, orPolywater, liquid water generally formed by condensation of water vapour in tiny glass or fused-quartz capillaries and with properties very different from those well established for ordinary water; e.g., lower vapour pressure, lower freezing temperature, higher density and viscosity, higher thermal stability, and different infrared and Raman spectra. For a few years after the announcement of the discovery of the substance (1968) by a group of Soviet scientists, many investigators held the view that the substance was a new form of water, possibly a polymer. In the 1970s thorough study established that anomalous water is ordinary water containing ionic contaminants that cause it to have the unusual properties.


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chemical element

oxygen (O), nonmetallic chemical element of Group 16 (VIa, or the oxygen group) of the periodic table. Oxygen is a colourless, odourless, tasteless gas essential to living organisms, being taken up by animals, which convert it to carbon dioxide; plants, in turn, utilize carbon dioxide as a source of carbon and return the oxygen to the atmosphere. Oxygen forms compounds by reaction with practically any other element, as well as by reactions that displace elements from their combinations with each other; in many cases, these processes are accompanied by the evolution of heat and light and in such cases are called combustions. Its most important compound is water.

Element Properties
atomic number 8
atomic weight 15.9994
melting point −218.4 °C (−361.1 °F)
boiling point −183.0 °C (−297.4 °F)
density (1 atm, 0 °C) 1.429 g/litre
oxidation states −1, −2, +2 (in compounds with fluorine)
electron config. 1s22s22p4


Oxygen was discovered about 1772 by a Swedish chemist, Carl Wilhelm Scheele, who obtained it by heating potassium nitrate, mercuric oxide, and many other substances. An English chemist, Joseph Priestley, independently discovered oxygen in 1774 by the thermal decomposition of mercuric oxide and published his findings the same year, three years before Scheele published. In 1775–80, French chemist Antoine-Laurent Lavoisier, with remarkable insight, interpreted the role of oxygen in respiration as well as combustion, discarding the phlogiston theory, which had been accepted up to that time; he noted its tendency to form acids by combining with many different substances and accordingly named the element oxygen (oxygène) from the Greek words for “acid former.”

Occurrence and properties

At 46 percent of the mass, oxygen is the most plentiful element in Earth’s crust. The proportion of oxygen by volume in the atmosphere is 21 percent and by weight in seawater is 89 percent. In rocks, it is combined with metals and nonmetals in the form of oxides that are acidic (such as those of sulfur, carbon, aluminum, and phosphorus) or basic (such as those of calcium, magnesium, and iron) and as saltlike compounds that may be regarded as formed from the acidic and basic oxides, as sulfates, carbonates, silicates, aluminates, and phosphates. Plentiful as they are, these solid compounds are not useful as sources of oxygen, because separation of the element from its tight combinations with the metal atoms is too expensive.

Below −183 °C (−297 °F), oxygen is a pale blue liquid; it becomes solid at about −218 °C (−361 °F). Pure oxygen is 1.1 times heavier than air.

During respiration, animals and some bacteria take oxygen from the atmosphere and return to it carbon dioxide, whereas by photosynthesis, green plants assimilate carbon dioxide in the presence of sunlight and evolve free oxygen. Almost all the free oxygen in the atmosphere is due to photosynthesis. About 3 parts of oxygen by volume dissolve in 100 parts of fresh water at 20 °C (68 °F), slightly less in seawater. Dissolved oxygen is essential for the respiration of fish and other marine life.

Natural oxygen is a mixture of three stable isotopes: oxygen-16 (99.759 percent), oxygen-17 (0.037 percent), and oxygen-18 (0.204 percent). Several artificially prepared radioactive isotopes are known. The longest-lived, oxygen-15 (124-second half-life), has been used to study respiration in mammals.


Oxygen has two allotropic forms, diatomic (O2) and triatomic (O3, ozone). The properties of the diatomic form suggest that six electrons bond the atoms and two electrons remain unpaired, accounting for the paramagnetism of oxygen. The three atoms in the ozone molecule do not lie along a straight line.

Ozone may be produced from oxygen according to the equation:Chemical equation.

The process, as written, is endothermic (energy must be provided to make it proceed); conversion of ozone back into diatomic oxygen is promoted by the presence of transition metals or their oxides. Pure oxygen is partly transformed into ozone by a silent electrical discharge; the reaction is also brought about by absorption of ultraviolet light of wavelengths around 250 nanometres (nm, the nanometre, equal to 10−9 metre); occurrence of this process in the upper atmosphere removes radiation that would be harmful to life on the surface of the Earth. The pungent odour of ozone is noticeable in confined areas in which there is sparking of electrical equipment, as in generator rooms. Ozone is light blue; its density is 1.658 times that of air, and it has a boiling point of −112 °C (−170 °F) at atmospheric pressure.

Ozone is a powerful oxidizing agent, capable of converting sulfur dioxide to sulfur trioxide, sulfides to sulfates, iodides to iodine (providing an analytical method for its estimation), and many organic compounds to oxygenated derivatives such as aldehydes and acids. The conversion by ozone of hydrocarbons from automotive exhaust gases to these acids and aldehydes contributes to the irritating nature of smog. Commercially, ozone has been used as a chemical reagent, as a disinfectant, in sewage treatment, water purification, and bleaching textiles.

Preparative methods

Production methods chosen for oxygen depend upon the quantity of the element desired. Laboratory procedures include the following:

1. Thermal decomposition of certain salts, such as potassium chlorate or potassium nitrate:Chemical equations.

The decomposition of potassium chlorate is catalyzed by oxides of transition metals; manganese dioxide (pyrolusite, MnO2) is frequently used. The temperature necessary to effect the evolution of oxygen is reduced from 400 °C to 250 °C by the catalyst.

2. Thermal decomposition of oxides of heavy metals:Chemical equations.

Scheele and Priestley used mercury(II) oxide in their preparations of oxygen.

3. Thermal decomposition of metal peroxides or of hydrogen peroxide:Chemical equations.

An early commercial procedure for isolating oxygen from the atmosphere or for manufacture of hydrogen peroxide depended on the formation of barium peroxide from the oxide as shown in the equations.

4. Electrolysis of water containing small proportions of salts or acids to allow conduction of the electric current:Chemical equation.

Commercial production and use

When required in tonnage quantities, oxygen is prepared by the fractional distillation of liquid air. Of the main components of air, oxygen has the highest boiling point and therefore is less volatile than nitrogen and argon. The process takes advantage of the fact that when a compressed gas is allowed to expand, it cools. Major steps in the operation include the following: (1) Air is filtered to remove particulates; (2) moisture and carbon dioxide are removed by absorption in alkali; (3) the air is compressed and the heat of compression removed by ordinary cooling procedures; (4) the compressed and cooled air is passed into coils contained in a chamber; (5) a portion of the compressed air (at about 200 atmospheres pressure) is allowed to expand in the chamber, cooling the coils; (6) the expanded gas is returned to the compressor with multiple subsequent expansion and compression steps resulting finally in liquefaction of the compressed air at a temperature of −196 °C; (7) the liquid air is allowed to warm to distill first the light rare gases, then the nitrogen, leaving liquid oxygen. Multiple fractionations will produce a product pure enough (99.5 percent) for most industrial purposes.

The steel industry is the largest consumer of pure oxygen in “blowing” high carbon steel—that is, volatilizing carbon dioxide and other nonmetal impurities in a more rapid and more easily controlled process than if air were used. The treatment of sewage by oxygen holds promise for more efficient treatment of liquid effluents than other chemical processes. Incineration of wastes in closed systems using pure oxygen has become important. The so-called LOX of rocket oxidizer fuels is liquid oxygen; the consumption of LOX depends upon the activity of space programs. Pure oxygen is used in submarines and diving bells.

Commercial oxygen or oxygen-enriched air has replaced ordinary air in the chemical industry for the manufacture of such oxidation-controlled chemicals as acetylene, ethylene oxide, and methanol. Medical applications of oxygen include use in oxygen tents, inhalators, and pediatric incubators. Oxygen-enriched gaseous anesthetics ensure life support during general anesthesia. Oxygen is significant in a number of industries that use kilns.

Chemical properties and reactions

The large values of the electronegativity and the electron affinity of oxygen are typical of elements that show only nonmetallic behaviour. In all of its compounds, oxygen assumes a negative oxidation state as is expected from the two half-filled outer orbitals. When these orbitals are filled by electron transfer, the oxide ion O2− is created. In peroxides (species containing the ion O22−) it is assumed that each oxygen has a charge of −1. This property of accepting electrons by complete or partial transfer defines an oxidizing agent. When such an agent reacts with an electron-donating substance, its own oxidation state is lowered. The change (lowering), from the zero to the −2 state in the case of oxygen, is called a reduction. Oxygen may be thought of as the “original” oxidizing agent, the nomenclature used to describe oxidation and reduction being based upon this behaviour typical of oxygen.

As described in the section on allotropy, oxygen forms the diatomic species, O2, under normal conditions and, as well, the triatomic species ozone, O3. There is some evidence for a very unstable tetratomic species, O4. In the molecular diatomic form there are two unpaired electrons that lie in antibonding orbitals. The paramagnetic behaviour of oxygen confirms the presence of such electrons.

The intense reactivity of ozone is sometimes explained by suggesting that one of the three oxygen atoms is in an “atomic” state; on reacting, this atom is dissociated from the O3 molecule, leaving molecular oxygen.

The molecular species, O2, is not especially reactive at normal (ambient) temperatures and pressures. The atomic species, O, is far more reactive. The energy of dissociation (O2 → 2O) is large at 117.2 kilocalories per mole.

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


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sulfur oxide

sulfur oxide, any of several compounds of sulfur and oxygen, the most important of which are sulfur dioxide (SO2) and sulfur trioxide (SO3), both of which are manufactured in huge quantities in intermediate steps of sulfuric acid manufacture. The dioxide is the acid anhydride (a compound that combines with water to form an acid) of sulfurous acid; the trioxide is the acid anhydride of sulfuric acid.

Sulfur dioxide is a heavy, colourless, poisonous gas with a pungent, irritating odour familiar as the smell of a just-struck match. Occurring in nature in volcanic gases and in solution in the waters of some warm springs, sulfur dioxide usually is prepared industrially by the burning in air or oxygen of sulfur or such compounds of sulfur as iron pyrite or copper pyrite. Large quantities of sulfur dioxide are formed in the combustion of sulfur-containing fuels; in the second half of the 20th century, measures to control atmospheric pollution by this compound were widely adopted. In the laboratory the gas may be prepared by reducing sulfuric acid (H2SO4) to sulfurous acid (H2SO3), which decomposes into water and sulfur dioxide, or by treating sulfites (salts of sulfurous acid) with strong acids, such as hydrochloric acid, again forming sulfurous acid.

Sulfur dioxide can be liquefied under moderate pressures at room temperatures; the liquid freezes at -73° C (-99.4° F) and boils at -10° C (+14° F) under atmospheric pressure. Although its chief uses are in the preparation of sulfuric acid, sulfur trioxide, and sulfites, sulfur dioxide also is used as a disinfectant, a refrigerant, a bleach, and a food preservative, especially in dried fruits.

Sulfur trioxide is a colourless compound that exists at room temperature either as a volatile liquid or in any of three allotropic solid forms. The liquid boils at 44.6° C (112° F) and solidifies at 16.83° C (62° F); the most stable of the solid forms melts at 62° C (144° F). Formed by the reaction of sulfur dioxide and oxygen in the presence of catalysts, sulfur trioxide fumes vigorously in contact with moist air and dissolves in water, liberating much heat and forming sulfuric acid. Solutions of the trioxide in sulfuric acid are called fuming sulfuric acid, or oleum. Like sulfuric acid, sulfur trioxide is a very powerful dehydrating agent, is very corrosive, and is very reactive chemically.

Other oxides of sulfur include the monoxide (SO), sesquioxide (S2O3), heptoxide (S2O7), and tetroxide (SO4). The monoxide is formed as an unstable colourless gas by an electric discharge in a mixture of sulfur dioxide and sulfur vapour at low pressure; upon cooling, it condenses to an orange-red solid that decomposes slowly to sulfur and sulfur dioxide. The sesquioxide, formed by dissolving sulfur in liquid sulfur trioxide, is a blue-green solid stable only below 15° C (59° F). The heptoxide and the tetroxide, unstable compounds that melt at about 0° C (32° F), are formed by an electric discharge in a mixture of sulfur dioxide or trioxide and oxygen.


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silane, also called Silicon Hydride, any of a series of covalently bonded compounds containing only the elements silicon and hydrogen, having the general formula SinH2n + 2, in which n equals 1, 2, 3, and so on. The silanes are structural analogues of the saturated hydrocarbons (alkanes) but are much less stable. The term silane is extended to include compounds in which any or all of the hydrogen atoms have been replaced by other atoms or groups of atoms, as in tetrachlorosilane, SiCl4.

Silanes have been prepared by the reaction of magnesium silicide (Mg2Si) with acids or by the reduction of silicon chlorides with lithium aluminum hydride. All the silanes burn or explode upon contact with air, and they are decomposed by alkaline solutions with formation of hydrogen and hydrous silica. Upon heating, the silanes decompose into hydrogen and silicon; they react with the halogens or hydrogen halides to form halogenated silanes, and with olefins to form alkylsilanes.

The simplest silane, monosilane (SiH4), is also the stablest; it is a colourless gas that liquefies at -112° C (-170° F) and freezes at -185° C (-301° F). It decomposes slowly at 250° C (482° F), rapidly at 500° C (932° F).

The instability of the silanes results from the reactivity of the silicon-hydrogen bond; derivatives in which all the hydrogen atoms have been replaced by organic groups, such as tetramethylsilane, Si(CH3)4, resemble the saturated hydrocarbons. The compound dimethyldichlorosilane, (CH3)2SiCl2, is important as the starting material for the dimethylpolysiloxanes, members of the silicone family of polymers. Chlorotrimethylsilane and vinyltrichlorosilane are used to impart water repellency to numerous materials such as cloth, paper, and glass.


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nonmetal, in physics, a substance having a finite activation energy (band gap) for electron conduction. This means that nonmetals display low (insulators) to moderate (semiconductors) bulk electrical conductivities, which increase with increasing temperature, and are subject to dielectric breakdown at high voltages and temperatures. Like metals, nonmetals may occur in the solid, liquid, or gaseous state. However, unlike metals, nonmetals display a wide range of both mechanical and optical properties, ranging from brittle to plastic and from transparent to opaque.

From a chemical point of view, nonmetals may be divided into two classes: 1) covalent materials, which contain atoms having small sizes, high electronegativities, low valence vacancy to electron ratios, and a pronounced tendency to form negative ions in chemical reactions and negative oxidation states in their compounds; 2) ionic materials, which contain both small and large atoms. Ions may be formed by adding electrons to (small, electronegative atoms) or by extracting electrons from (large, electropositive) atoms. In ionic materials, nonmetals exist either as monatomic anions (e. g., F-in NaF) or as constituents of polyatomic anions (e.g., N and O in the NO3-`s in NaNO3). When in the form of simple elemental substances, about 25 or 22% of the known elements form nonmetals at normal temperatures and pressures, including all of the elements in the S-block of the periodic table and approximately 58% of those in the P-block.


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Jean André Deluc

Swiss-British geologist and meteorologist

Jean André Deluc, (born Feb. 8, 1727, Geneva, Switz.—died Nov. 7, 1817, Windsor, Berkshire, Eng.), Swiss-born British geologist and meteorologist whose theoretical work was influential on 19th-century writing about meteorology.

Deluc was educated in mathematics and the natural sciences. He engaged in business, and on his business travels around Europe he collected mineral and plant specimens.

Deluc suffered business reverses in 1773 and left Geneva for England. He devoted himself to his scientific interests and was made a fellow of the Royal Society. He became a reader to Queen Charlotte, which allowed him much time and the means to travel on the European continent. He studied the effects of heat and pressure on the mercury barometer and, as a pioneer in scientific mountaineering, published the first correct rules for using the barometer to find the heights of mountains.

Deluc discovered that water attains its maximum density at 39° F (4° C), and he developed the theory that the quantity of water vapour in any given space is independent of the density of the air in which it is diffused. As an amateur physicist, he constructed an “electric column” of zinc and silvered paper similar to the galvanic pile of Alessandro Volta, which was in vogue for electrical experimentation for some time. Deluc’s keenest personal interest was his quest to reconcile the Creation story of Genesis with the evidence of geology, and to this end he interpreted each day of the Creation as an epoch.


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Roman religion

Egeria, in Roman religion, a water spirit worshiped in connection with Diana at Aricia and also with the Camenae in their grove outside the Porta Capena at Rome. Like Diana, she was a protectress of pregnant women and, like the Camenae, was considered prophetic. Traditionally she was the wife, or mistress, and adviser of King Numa Pompilius, who established the grove at Rome and consorted with her there.


Now commonly used as body armor, Kevlar fabric was invented in an attempt to make lightweight tires to increase car gas mileage.

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