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Water (molecule)

Water (molecule)
Water (H2O) Molecular shape Dipole moment Hazards Main hazards None (see also Dihydrogen monoxide hoax) Hydrogen sulfide Hydrogen selenide Hydrogen telluride acetone methanol water vapor ice heavy water bent 1.85 D

IUPAC name Water Other names Dihydrogen monoxide Oxidane Hydroxylic Acid Hydrogen Hydroxide [7732-18-5] 15377 ZC0110000 H2 O 18.01528(33) g/mol white solid or almost colourless, transparent, with a slight hint of blue, crystalline solid or liquid [1] 1000 kg/m3, liquid (4 °C) 917 kg/m3, solid 0 °C, 32 °F (273.15 K)[2] 100 °C, 212 °F (373.15 K)[2] 15.74 ~35-36 15.74 1.3330 0.001 Pa s at 20 °C Hexagonal See ice

Related compounds Other cations

Related solvents Related compounds

Identifiers CAS number ChEBI RTECS number Properties Molecular formula Molar mass Appearance

Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox references

Density

Melting point Boiling point Acidity (pKa) Basicity (pKb) Refractive index (nD) Viscosity Structure Crystal structure

A pool lined with white tiles shows the intrinsic light blue color of water. Water (H2O, HOH) is the most abundant molecule on Earth’s surface, constituting about 75% of the Earth’s surface in liquid, solid, and gaseous states. It is in dynamic equilibrium between the liquid and gas states at standard temperature and pressure. At room temperature, it is a nearly colorless (with a hint of blue), tasteless, and odorless liquid. Many substances dissolve in water and it is commonly referred to as the universal solvent. Because of this, water in nature and in use is rarely pure and some of its properties may vary slightly from those of the

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pure substance. However, there are many compounds that are essentially, if not completely, insoluble in water. Water is the only common substance found naturally in all three common states of matter—for other substances, see Chemical properties. Water is essential for life on Earth. Water usually makes up 55% to 78% of the human body.[3]

Water (molecule)
hydrogen to produce gases under standard conditions. The reason that water forms a liquid is that oxygen is more electronegative than all of these elements with the exception of fluorine. Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom. The presence of a charge on each of these atoms gives each water molecule a net dipole moment. Electrical attraction between water molecules due to this dipole pulls individual molecules closer together, making it more difficult to separate the molecules and therefore raising the boiling point. This attraction is known as hydrogen bonding. The molecules of water are constantly moving in relation to each other, and the hydrogen bonds are continually breaking and reforming at the timescales faster than 200 femtoseconds.[6] However, this bond is strong enough to create many of the peculiar properties of water described in this article, such as the ones that make it integral to life. Water can be described as a polar liquid that slightly dissociates disproportionately into the hydronium ion (H3O+(aq)) and an associated hydroxide ion (OH−(aq)). 2 H2O (l) H3O+(aq) + OH−(aq) The dissociation constant for this dissociation is commonly symbolized as Kw and has a value of about 10-14 at 25°C; see "Water (data page)" and "Self-ionization of water" for more information.

Forms of water
Water can take many forms. The solid state of water is known as ice; the gaseous state is known as water vapor (or steam), and the common liquid phase is generally understood when simply referring to water. Above a certain critical temperature and pressure (647 K and 22.064 MPa), water molecules assume a supercritical condition, in which liquid-like clusters float within a vapor-like phase. There are many different crystalline and amorphous forms of ice; see ice for a complete listing. In natural water, almost all of the hydrogen atoms are of the isotope protium, 1H. Heavy water is water in which the hydrogen is replaced by its heavier isotope, deuterium, 2H. It is chemically almost identical to normal water. Heavy water is used in the nuclear reactor industry to moderate (slow down) neutrons. By contrast, the term light water designates water containing the protium isotope, in situation when such distinction is needed. An example is the term light water reactor to emphasize that the reactor type uses light water.

Water, ice and vapor
Heat capacity and heats of vaporization and fusion
Water has the second highest specific heat capacity of all known substances, after ammonia, as well as a high heat of vaporization (40.65 kJ mol−1), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth’s climate by buffering large fluctuations in temperature. The specific enthalpy of fusion of water is 333.55 kJ kg−1 at 0 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting upon the ice of glaciers and drift ice. Before the advent of mechanical refrigeration, ice was in common use to retard food spoilage.

Physics and chemistry
Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. [4] Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water vapor is essentially invisible as a gas.[5] Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the oxygen family in the periodic table, which are gases such as hydrogen sulfide. Also the elements surrounding oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, all combine with

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Water (molecule)
freezing. It can remain in a fluid state down to its crystal homogeneous nucleation at almost 231 K (−42 °C) [8]. Water also expands significantly as the temperature increases. Its density decreases by 4% from its highest value when approaching the boiling point. These properties of water have important consequences in its role in the ecosystem of Earth. Water of a temperature of 4°C will always accumulate at the bottom of fresh water lakes, irrespective of the temperature in the atmosphere. Since water and ice are poor conductors of heat (good insulators) it is unlikely that sufficiently deep lakes will freeze completely, unless stirred by strong currents that would intermix cooler and warmer water and facility more rapid cool-off. In warming weather, chunks of ice float, rather than sink to the bottom where they might melt extremely slowly. These phenomena thus may preserve aquatic life.

Density of water and ice
Temp (°C) +100 +80 +60 +40 +30 +25 +22 +20 +15 +10 +4 0 −10 −20 −30 Density (kg/m³) 958.4 971.8 983.2 992.2 995.6502 997.0479 997.7735 998.2071 999.1026 999.7026 999.9720 999.8395 998.117 993.547 983.854

The density of water (in kg/m3) at various temperatures in degrees Celsius [7] The values below 0 °C refer to supercooled water.

Density of saltwater and ice
The density of water is dependent on the dissolved salt content as well as the temperature of the water. Ice still floats in the oceans, otherwise they would freeze from the bottom up. However, the salt content of oceans lowers the freezing point by about 2 °C and lowers the temperature of the density maximum of water to the freezing point. That is why, in ocean water, the downward convection of colder water is not blocked by an expansion of water as it becomes colder near the freezing point. The oceans’ cold water near the freezing point continues to sink. For this reason, any creature attempting to survive at the bottom of such cold water as the Arctic Ocean generally lives in water that is 4 °C colder than the temperature at the bottom of frozen-over fresh water lakes and rivers in the winter. As the surface of salt water begins to freeze (at −1.9 °C for normal salinity seawater, 3.5%) the ice that forms is essentially salt free with a density approximately equal to that of freshwater ice. This ice floats on the surface and the salt that is "frozen out" adds to the salinity and density of the seawater just below it, in a process known as brine rejection. This denser saltwater sinks by convection and the replacing seawater is subject to the same process. This provides essentially freshwater ice at −1.9 °C on the surface. The increased density of the seawater beneath

Water - Density and Specific Weight The density of water is dependent on its temperature, but the relation is not linear and is not even monotonic (see right-hand table). The solid form of most substances is more dense than the liquid phase; thus, a block of the solid will sink in a tub of the liquid. But, by contrast, a block of common ice floats in a tub of water because ice is less dense than liquid water. When cooled from room temperature liquid water becomes increasingly denser, just like other substances. But at approximately 4 °C, water reaches its maximum density. As it is cooled further under ambient conditions, it expands to become less dense. Upon freezing, the density of ice decreases by about 9%. The physical reason of this is related to the crystal structure of ordinary ice, known as hexagonal ice Ih. Not all forms of ice are less dense than liquid water however, high density amorphous ice (HDA) and very high density amorphous ice (VHDA), for example, are both denser than pure liquid water. The melting point of water is 0 °C (32 °F, 273 K), however, liquid water can be supercooled well below that temperature without

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the forming ice causes it to sink towards the bottom. On a large scale, the process of brine rejection and sinking cold salty water results in ocean currents forming to transport such water away from the pole. One potential consequence of global warming is that the loss of Arctic ice could result in the loss of these currents as well, which could have unforeseeable consequences on near and distant climates.

Water (molecule)
saturated or 100% relative humidity, when the vapor pressure of water in the air is at the equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool enough) will fail to lose mass through evaporation when exposed to saturated air. Because the amount of water vapor in air is small, relative humidity, the ratio of the partial pressure due to the water vapor to the saturated partial vapor pressure, is much more useful. Water vapor pressure above 100% relative humidity is called super-saturated and can occur if air is rapidly cooled, say by rising suddenly in an updraft.[10]

Miscibility and condensation

Vapor pressure
[11]

Compressibility
The compressibility of water is a function of pressure and temperature. At 0 °C in the limit of zero pressure the compressibility is 5.1×10-5 bar−1.[12] In the zero pressure limit the compressibility reaches a minimum of 4.4×10-5 bar−1 around 45 °C before increasing again with increasing temperature. As the pressure is increased the compressibility decreases, being 3.9×10-5 bar−1 at 0 °C and 1000 bar. The bulk modulus of water is 2.2×109 Pa.[13] The low compressibility of non-gases, and of water in particular, leads to them often being assumed as incompressible. The low compressibility of water means that even in the deep oceans at 4000 m depth, where pressures are 4×107 Pa, there is only a 1.8% decrease in volume.[13]

Red line shows saturation Water is miscible with many liquids, for example ethanol in all proportions, forming a single homogeneous liquid. On the other hand water and most oils are immiscible usually forming layers according to increasing density from the top. As a gas, water vapor is completely miscible with air. On the other hand the maximum water vapor pressure that is thermodynamically stable with the liquid (or solid) at a given temperature is relatively low compared with total atmospheric pressure. For example, if the vapor partial pressure[9] is 2% of atmospheric pressure and the air is cooled from 25 °C, starting at about 22 °C water will start to condense, defining the dew point, and creating fog or dew. The reverse process accounts for the fog burning off in the morning. If one raises the humidity at room temperature, say by running a hot shower or a bath, and the temperature stays about the same, the vapor soon reaches the pressure for phase change, and condenses out as steam. A gas in this context is referred to as

Triple point
The temperature and pressure at which solid, liquid, and gaseous water coexist in equilibrium is called the triple point of water. This point is used to define the units of temperature (the kelvin, the SI unit of thermodynamic temperature and, indirectly, the degree Celsius and even the degree Fahrenheit). As a consequence, water’s triple point temperature is a prescribed value rather than a measured quantity. The triple point is at a temperature of 273.16 K (0.01 °C) by convention, and at a pressure of 611.73 Pa. This pressure is quite low, about 1/166 of the normal sea level barometric pressure of 101,325 Pa. The atmospheric surface pressure on planet Mars is remarkably close to the triple point pressure, and the zero-elevation or "sea level" of Mars

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Temperature (°C) 0 5 10 12 14 16 17 18 19 20 21 22 23 24 25 Pressure (torr) 4.58 6.54 9.21 10.52 11.99 13.63 14.53 15.48 16.48 17.54 18.65 19.83 21.07 22.38 23.76 The various triple points of water[14] Phases in stable equilibrium liquid water, ice Ih, and water vapour liquid water, ice Ih, and ice III liquid water, ice III, and ice V liquid water, ice V, and ice VI ice Ih, Ice II, and ice III ice II, ice III, and ice V ice II, ice V, and ice VI Pressure 611.73 Pa 209.9 MPa 350.1 MPa 632.4 MPa 213 MPa 344 MPa 626 MPa

Water (molecule)
Pressure (mbar) 6.11 8.72 12.28 14.03 15.99 18.17 19.37 20.64 21.97 23.38 24.86 26.44 28.09 29.84 31.68

Temperature 273.16 K (0.01 °C) 251 K (-22 °C) -17.0 °C 0.16 °C -35 °C -24 °C -70 °C

is defined by the height at which the atmospheric pressure corresponds to the triple point of water. Although it is commonly named as "the triple point of water", the stable combination of liquid water, ice I, and water vapour is but one of several triple points on the phase diagram of water. Gustav Heinrich Johann Apollon Tammann in Göttingen produced data on several other triple points in the early 20th century. Kamb and others documented further triple points in the 1960s.[14][15][16] water phase diagram: Y-axis = Pressure in pascal (10n), X-axis = Temperature in kelvin, S = Solid, L = Liquid, V = Vapour, CP = Critical Point, TP = Triple point of water

Electrical properties
Electrical conductivity
Pure water containing no ions is an excellent insulator, but not even "deionized" water is completely free of ions. Water undergoes auto-ionisation in the liquid state. Further,

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because water is such a good solvent, it almost always has some solute dissolved in it, most frequently a salt. If water has even a tiny amount of such an impurity, then it can conduct electricity readily, as impurities such as salt separate into free ions in aqueous solution by which an electric current can flow. It is known that the theoretical maximum electrical resistivity for water is approximately 182 kΩ·m at 25 °C. This figure agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultrapure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in ultra-pure water begins to noticeably lower its resistivity level by up to several kiloohm-meters (or hundreds of nanosiemens per meter). The low electrical conductivity of water increases significantly upon solvation of a small amount of ionic material, such as hydrogen chloride or any salt. Thus the risks of electrocution are much greater in water with impurities. It is worth noting, however, that the risk of electrocution decreases when the impurities increase to the point at which the water itself is a better conductor than the human body. For example, the risks of electrocution in sea water may be lower than in fresh water, as the sea has a much higher level of impurities, particularly common salt. The main current path will seek the better conductor. Any electrical conductivity observable in water is the result of ions of mineral salts and carbon dioxide dissolved in it. Carbon dioxide forms carbonate ions in Water. Water self-ionizes in which two water molecules form one hydroxide anion and one hydronium cation, but not enough to carry enough electric current to do any work or harm for most operations. In pure water, sensitive equipment can detect a very slight electrical conductivity of 0.055 µS/cm at 25 °C. Water can also be electrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is a very slow process, as very little current is conducted. While electrons are the primary charge carriers in water (and metals), in ice the primary charge carriers are protons (see proton conductor).

Water (molecule)

Electrolysis
Water can be split into its constituent elements, hydrogen and oxygen, by passing an electric current through it. This process is called electrolysis. Water molecules naturally dissociate into H+ and OH− ions, which are pulled toward the cathode and anode, respectively. At the cathode, two H+ ions pick up electrons and form H2 gas. At the anode, four OH− ions combine and release O2 gas, molecular water, and four electrons. The gases produced bubble to the surface, where they can be collected. The standard potential of the water electrolysis cell is 1.23 V at 25 °C.

Dipolar properties

model of hydrogen bonds between molecules of water An important feature of water is its polar nature. The water molecule forms an angle, with hydrogen atoms at the tips and oxygen at the vertex. Since oxygen has a higher electronegativity than hydrogen, the side of the molecule with the oxygen atom has a partial negative charge. An object with such a charge difference is called a dipole. The charge differences cause water molecules to be attracted to each other (the relatively positive areas being attracted to the relatively negative areas) and to other polar molecules. This attraction contributes to hydrogen bonding, and explains many of the properties of water, such as solvent action. Water’s dipolar nature can be demonstrated by holding an electrically charged object (such as a comb

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after combing) near a small falling stream of water (e.g., from a faucet), causing the stream to be attracted towards the charged object. Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for a number of water’s physical properties. One such property is its relatively high melting and boiling point temperatures; more energy is required to break the hydrogen bonds between molecules. The similar compound hydrogen sulfide (H2S), which has much weaker hydrogen bonding, is a gas at room temperature even though it has twice the molecular mass of water. The extra bonding between water molecules also gives liquid water a large specific heat capacity. This high heat capacity makes water a good heat storage medium (coolant) and heat shield. Hydrogen bonding also gives water its unusual behavior when freezing. When cooled to near freezing point, the presence of hydrogen bonds means that the molecules, as they rearrange to minimize their energy, form the hexagonal crystal structure of ice that is actually of lower density: hence the solid form, ice, can float in water. In other words, water expands as it freezes, whereas almost all other materials shrink on solidification. Due to the temperature dependent properties of hydrogen bonding, Water has its highest density at approx. 4 °C. In freezing winter weather, water of this temperature accumulates at the bottom of lakes before total freezing may occur. An interesting consequence of the solid having a lower density than the liquid is that ice will melt if sufficient pressure is applied. With increasing pressure the melting point temperature drops and when the melting point temperature is lower than the ambient temperature the ice begins to melt. A significant increase of pressure is required to lower the melting point temperature —the pressure exerted by an ice skater on the ice would only reduce the melting point by approximately 0.09 °C (0.16 °F).

Water (molecule)

Dew drops adhering to a spider web between glass and water molecules (adhesive forces) are stronger than the cohesive forces. In biological cells and organelles, water is in contact with membrane and protein surfaces that are hydrophilic; that is, surfaces that have a strong attraction to water. Irving Langmuir observed a strong repulsive force between hydrophilic surfaces. To dehydrate hydrophilic surfaces—to remove the strongly held layers of water of hydration—requires doing substantial work against these forces, called hydration forces. These forces are very large but decrease rapidly over a nanometer or less. Their importance in biology has been extensively studied by V. Adrian Parsegian of the National Institute of Health.[17] They are particularly important when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.

Surface tension
Water has a high surface tension of 72.8 mN/ m, caused by the strong cohesion between water molecules, the highest of the nonmetallic liquids. This can be seen when small quantities of water are placed onto a sorption-free (non-adsorbent and non-absorbent) surface, such as polythene or Teflon, and the water stays together as drops. Just as significantly, air trapped in surface disturbances forms bubbles, which sometimes last long enough to transfer gas molecules to the water. Another surface tension effect is capillary waves, which are the surface ripples that form around the impacts of drops on water surfaces, and sometimes occur with strong subsurface currents flowing to the water surface. The apparent elasticity caused by surface tension drives the waves.

Adhesion
Water sticks to itself (cohesion) because it is polar. Water also has high adhesion properties because of its polar nature. On extremely clean/smooth glass the water may form a thin film because the molecular forces

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Water (molecule)
pockets and rendering the xylem water transport inoperative.

Water as a solvent

This paper clip is under the water level, which has risen gently and smoothly. Surface tension prevents the paper clip from submerging and from overflowing the blue glass.

Capillary action
Due to an interplay of the forces of adhesion and surface tension, water exhibits capillary action whereby water rises into a narrow tube against the force of gravity. Water adheres to the inside wall of the tube and surface tension tends to straighten the surface causing a surface rise and more water is pulled up through cohesion. The process continues as the water flows up the tube until there is enough water such that gravity balances the adhesive force. Surface tension and capillary action are important in biology. For example, when water is carried through xylem up stems in plants, the strong intermolecular attractions (cohesion) hold the water column together and strong adhesive properties maintain the water attachment to the xylem and prevent tension rupture caused by transpiration pull. Other liquids with lower surface tension and poorer capillary action would have a higher tendency to disrupt the column, forming air

High concentrations of dissolved lime make the water of Havasu Falls turn turquoise. Water is also a good solvent due to its polarity. Substances that will mix well and dissolve in water (e.g. salts) are known as hydrophilic ("water-loving") substances, while those that do not mix well with water (e.g. fats and oils), are known as hydrophobic ("water-fearing") substances. The ability of a substance to dissolve in water is determined by whether or not the substance can match or better the strong attractive forces that water molecules generate between other water molecules. If a substance has properties that do not allow it to overcome these strong intermolecular forces, the molecules are "pushed out" from the water, and do not dissolve. Contrary to the common misconception, water and hydrophobic substances do not "repel", and the hydration of a hydrophobic surface is energetically, but not entropically, favorable. When an ionic or polar compound enters water, it is surrounded by water molecules

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(Hydration). The relatively small size of water molecules typically allows many water molecules to surround one molecule of solute. The partially negative dipole ends of the water are attracted to positively charged components of the solute, and vice versa for the positive dipole ends. In general, ionic and polar substances such as acids, alcohols, and salts are relatively soluble in water, and nonpolar substances such as fats and oils are not. Nonpolar molecules stay together in water because it is energetically more favorable for the water molecules to hydrogen bond to each other than to engage in van der Waals interactions with nonpolar molecules. An example of an ionic solute is table salt; the sodium chloride, NaCl, separates into Na+ cations and Cl- anions, each being surrounded by water molecules. The ions are then easily transported away from their crystalline lattice into solution. An example of a nonionic solute is table sugar. The water dipoles make hydrogen bonds with the polar regions of the sugar molecule (OH groups) and allow it to be carried away into solution.

Water (molecule)
instance, it receives an H+ ion from HCl in the equilibrium: HCl + H2O ⇌ H3O+ + Cl− Here water is acting as a base, by receiving an H+ ion. In the reaction with ammonia, NH3, water donates an H+ ion, and is thus acting as an acid: NH3 + H2O ⇌ NH4+ + OH−

Acidity in nature
In theory, pure water has a pH of 7 at 298 K. In practice, pure water is very difficult to produce. Water left exposed to air for any length of time will rapidly dissolve carbon dioxide, forming a dilute solution of carbonic acid, with a limiting pH of about 5.7. As cloud droplets form in the atmosphere and as raindrops fall through the air minor amounts of CO2 are absorbed and thus most rain is slightly acidic. If high amounts of nitrogen and sulfur oxides are present in the air, they too will dissolve into the cloud and rain drops producing more serious acid rain problems.

Ligand chemistry
The water molecule can also be used as a ligand in transition metal complexes; one example is perrhenic acid, which forms when Re2O7 is exposed to water. It contains two water molecules coordinated to a rhenium atom.

Hydrogen bonding
A water molecule can form a maximum of four hydrogen bonds because it can accept two and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia, methanol form hydrogen bonds but they do not show anomalous behaviour of thermodynamic, kinetic or structural properties like those observed in water. The answer to the apparent difference between water and other hydrogen bonding liquids lies in the fact that apart from water none of the hydrogen bonding molecules can form four hydrogen bonds either due to an inability to donate/accept hydrogens or due to steric effects in bulky residues. In water local tetrahedral order due to the four hydrogen bonds gives rise to an open structure and a 3-dimensional bonding network, which exists in contrast to the closely packed structures of simple liquids. There is a great similarity between water and silica in their anomalous behaviour, even though one (water) is a liquid which has a hydrogen bonding network while the other (silica) has a covalent network with a very high melting point. One reason that water is well suited, and chosen, by life-forms, is that it exhibits its unique properties over a

Amphoteric nature
Chemically, water is amphoteric — i.e., it is able to act as either an acid or a base. Occasionally the term hydroxic acid is used when water acts as an acid in a chemical reaction. At a pH of 7 (neutral), the concentration of hydroxide ions (OH−) is equal to that of the hydronium (H3O+) or hydrogen (H+) ions. If the equilibrium is disturbed, the solution becomes acidic (higher concentration of hydronium ions) or basic (higher concentration of hydroxide ions). Water can act as either an acid or a base in reactions. According to the BrønstedLowry system, an acid is defined as a species which donates a proton (an H+ ion) in a reaction, and a base as one which receives a proton. When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid. For

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temperature regime that suits diverse biological processes, including hydration. It is believed that hydrogen bond in water is largely due to electrostatic forces and some amount of covalency. The partial covalent nature of hydrogen bond predicted by Linus Pauling in the 1930s is yet to be proven unambiguously by experiments and theoretical calculations.

Water (molecule)
occurs naturally in ordinary water in very low concentrations (~0.03%) and D2O in far lower amounts (0.000003%). Consumption of pure isolated D2O may affect biochemical processes - ingestion of large amounts impairs kidney and central nervous system function. However, very large amounts of heavy water must be consumed for any toxicity to be apparent, and smaller quantities can be consumed with no ill effects at all. Oxygen also has three stable isotopes, with 16O present in 99.76 %, 17O in 0.04% and 18O in 0.2% of water molecules.[21]

Ultra-short lived quantum effects
Although the molecular formula of water is generally considered to be a stable result in molecular science, recent work started in 1995 has shown that at certain time scales (attoseconds = 10-18s) in electron and neutron scattering experiments, water may behave more like H3/2O than H2O at the quantum level.[18] This result could have significant ramifications to the understanding of short-lived water chemistry and physics. The experiments show that when high-energy neutrons and electrons collide with water, they scatter in a way that indicates that they are affected by a ratio of only 1.5:1 for the hydrogen to oxygen ratio. In the scattering process, the particle probes cause bond breakage in water and the experiments afford an insight into ultra-fast reaction chemistry that can only be observed in highly resolved kinetic and dynamical systems.[19][20]

Transparency
Water is relatively transparent to visible light, near ultraviolet light, and far-red light, but it absorbs most ultraviolet light, infrared light, and microwaves. Most photoreceptors and photosynthetic pigments utilize the portion of the light spectrum that is transmitted well through water. Microwave ovens take advantage of water’s opacity to microwave radiation to heat the water inside of foods. The very weak onset of absorption in the red end of the visible spectrum lends water its intrinsic blue hue (see Color of water).

History
The properties of water have historically been used to define various temperature scales. Notably, the Kelvin, Celsius and Fahrenheit scales were, or currently are, defined by the freezing and boiling points of water. The less common scales of Delisle, Newton, Réaumur and Rømer were defined similarly. The triple point of water is a more commonly used standard point today.[22]

Heavy water and isotopologues
There are several isotopes of both hydrogen and oxygen, so several isotopologues of water are known. Hydrogen has three naturally occurring isotopes. The most common, making up more than 99.98% of the hydrogen in water, has 1 proton and 0 neutrons. A second isotope, deuterium (short form "D"), has 1 proton and 1 neutron. Deuterium oxide, D2O, is also known as heavy water and is used in nuclear reactors as a neutron moderator. The third isotope, tritium, has 1 proton and 2 neutrons, and is radioactive, with a half-life of 4500 days. T2O exists in nature only in tiny quantities, being produced primarily via cosmic ray-driven nuclear reactions in the atmosphere. D2O is stable, but differs from H2O in that it is denser - hence, "heavy water" - and in that several other physical properties are slightly different from those of common, Hydrogen-1 containing "light water". Water with one deuterium atom HDO

Use
The first scientific decomposition of water into hydrogen and oxygen, by electrolysis, was done in 1800 by William Nicholson, an English chemist. In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen (by volume). Gilbert Newton Lewis isolated the first sample of pure heavy water in 1933. Polywater was a hypothetical polymerized form of water that was the subject of much scientific controversy during the late 1960s. The consensus now is that it does not exist.

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Water (molecule)
• Water dimer • Water model

Systematic naming
The accepted IUPAC name of water is simply "water" (or its equivalent in a different language), although there are two other systematic names which can be used to describe the molecule. The simplest and best systematic name of water is hydrogen oxide. This is analogous to related compounds such as hydrogen peroxide, hydrogen sulfide, and deuterium oxide (heavy water). Another systematic name, oxidane, is accepted by IUPAC as a parent name for the systematic naming of oxygenbased substituent groups,[23] although even these commonly have other recommended names. For example, the name hydroxyl is recommended over oxidanyl for the –OH group. The name oxane is explicitly mentioned by the IUPAC as being unsuitable for this purpose, since it is already the name of a cyclic ether also known as tetrahydropyran in the Hantzsch-Widman system; similar compounds include dioxane and trioxane.

References
[1] Braun, Charles L.; Sergei N. Smirnov (1993). "Why is water blue?" (HTML). J. Chem. Educ. 70 (8): 612. http://www.dartmouth.edu/~etrnsfer/ water.htm. [2] ^ Vienna Standard Mean Ocean Water (VSMOW), used for calibration, melts at 273.1500089(10) K (0.000089(10) °C, and boils at 373.1339 K (99.9839 °C) [3] Re: What percentage of the human body is composed of water? Jeffrey Utz, M.D., The MadSci Network [4] Campbell, Neil A.; Brad Williamson; Robin J. Heyden (2006). Biology: Exploring Life. Boston, Massachusetts: Pearson Prentice Hall. ISBN 0-13-250882-6. http://www.phschool.com/ el_marketing.html. [5] Braun, Charles L.; Sergei N. Smirnov (1993). "Why is water blue?" (HTML). J. Chem. Educ. 70 (8): 612. http://www.dartmouth.edu/~etrnsfer/ water.htm. [6] Smith, Jared D. (2005). "Unified description of temperature-dependent hydrogen bond rearrangements in liquid water"". Proc. Natl. Acad. Sci 102 (40): 14171–14174. doi:10.1073/ pnas.0506899102. http://www.lbl.gov/ Science-Articles/Archive/sabl/2005/ October/Hydrogen-bonds-in-liquidwater.pdf. [7] Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Edn.). Boca Raton (FL):CRC Press. [8] P. G. Debenedetti, P. G., and Stanley, H. E.; "Supercooled and Glassy Water", Physics Today 56 (6), p. 40–46 (2003). [9] The pressure due to water vapor in the air is called the partial pressure (Dalton’s law) and it is directly proportional to concentration of water molecules in air (Boyle’s law). [10] Adiabatic cooling resulting from the ideal gas law. [11] Brown, Theodore L., H. Eugene LeMay, Jr., and Bruce E. Burston. Chemistry: The Central Science. 10th ed. Upper Saddle River, NJ: Pearson Education, Inc., 2006.

Systematic nomenclature
Dihydrogen monoxide or DHMO is an overly pedantic systematic covalent name of water. This term has been used in parodies of chemical research that call for this "lethal chemical" to be banned. While the name "dihydrogen monoxide" is technically correct other systematic names for water include hydroxic acid, hydroxylic acid, and hydrogen hydroxide (both acid and alkali names exist for water because it is amphoteric, able to react both as an acid or an alkali, depending on the strength of the acid or alkali it is reacted with) - none of these names are used widely outside of DHMO sites. That said, some material safety data sheets for water will list drowning as a hazard.[24][25]

See also
• • • • • • • Double distilled water Hydrodynamics Superheated water Vienna Standard Mean Ocean Water Viscosity of Water Water (data page) Water absorption of electromagnetic radiation • Water cluster

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From Wikipedia, the free encyclopedia
[12] Fine, R.A. and Millero, F.J. (1973). "Compressibility of water as a function of temperature and pressure". Journal of Chemical Physics 59 (10): 5529. doi:10.1063/1.1679903. [13] ^ R. Nave. "Bulk Elastic Properties". HyperPhysics. Georgia State University. http://hyperphysics.phy-astr.gsu.edu/ hbase/hph.html. Retrieved on 2007-10-26. [14] ^ Oliver Schlüter (2003-07-28) (PDF). Impact of High Pressure — Low Temperature Processes on Cellular Materials Related to Foods. Technischen Universität Berlin. http://edocs.tuberlin.de./diss/2003/schlueter_oliver.pdf. [15] Gustav Heinrich Johann Apollon Tammann (1925). The States Of Aggregation. Constable And Company Limited. [16] William Cudmore McCullagh Lewis and James Rice (1922). A System of Physical Chemistry. Longmans, Green and co.. [17] Physical Forces Organizing Biomolecules (PDF) [18] Phil Schewe, James Riordon, and Ben Stein (31 Jul 03). "A Water Molecule’s Chemical Formula is Really Not H2O". Physics News Update. http://www.aip.org/enews/physnews/ 2003/split/648-1.html. [19] C. A. Chatzidimitriou-Dreismann, T. Abdul Redah, R. M. F. Streffer and J. Mayers (1997). "Anomalous Deep Inelastic Neutron Scattering from Liquid H2O-D2O: Evidence of Nuclear Quantum Entanglement". Physical Review Letters 79 (15): 2839. doi:10.1103/ PhysRevLett.79.2839. [20] C. A. Chatzidimitriou-Dreismann, M. Vos, C. Kleiner and T. Abdul-Redah (2003). "Comparison of Electron and Neutron Compton Scattering from Entangled Protons in a Solid Polymer". Physical

Water (molecule)
Review Letters 91 (5): 057403–4. doi:10.1103/PhysRevLett.91.057403. [21] IAPWS (2001). "Guideline on the Use of Fundamental Physical Constants and Basic Constants of Water". http://www.iapws.org/relguide/ fundam.pdf. [22] A Brief History of Temperature Measurement [23] Leigh, G. J. et al. 1998. Principles of chemical nomenclature: a guide to IUPAC recommendations, p. 99. Blackwell Science Ltd, UK. ISBN 0-86542-685-6 [24] MSDS David Grays Distilled Water 060106.pdf, HEALTH EFFECTS INHALED: "...excessive inhalation may cause drowning." [25] MSDS for BATTERY WATER, SECTION VI - Health Hazard Data: "WATER MAY CAUSE DEATH BY DROWNING"

External links
• Release on the IAPWS Industrial Formulation 1997 for the Thermodynamic Properties of Water and Steam (fast computation speed) • Release on the IAPWS Formulation 1995 for the Thermodynamic Properties of Ordinary Water Substance for General and Scientific Use (simpler formulation) • Sigma Xi The Scientific Research Society, Year of Water 2008 • Stockholm International Water Institute (SIWI) • Explanation of the anomalous properties of water • Water phase diagrams • Water Absorption Spectrum • Calculation of vapor pressure, liquid density, dynamic liquid viscosity, surface tension of water

Retrieved from "http://en.wikipedia.org/wiki/Water_(molecule)" Categories: Wikipedia articles needing reorganization, Forms of water, Hydrides, Hydrogen compounds, Hydroxides, Oxides, Inorganic solvents, Water chemistry, Neutron moderators This page was last modified on 20 May 2009, at 17:24 (UTC). All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.) Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a U.S. registered 501(c)(3) taxdeductible nonprofit charity. Privacy policy About Wikipedia Disclaimers

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