View source for Water (molecule)

{{Chembox new
|   Name = Water (H<sub>2</sub>O)
|   ImageFile = Water-2D-labelled.png
|   ImageSize = 160px
|   ImageName = The water molecule has this basic geometric structure
|   ImageFile1 = Water molecule 3D.svg
|   ImageSize1 = 120px
|   ImageName1 = Water molecules have this structure.
|   IUPACName = Water
|   OtherNames = Aqua<br (Aq) />Hydrogen oxide<br />Hydrogen hydroxide<br />Hydrate<br />Oxidane<br />Hydroxic acid<br />[[Dihydrogen monoxide]]<br />Hydroxyl acid<br />Dihydrogen oxide<br />Hydrohydroxic acid<br />''μ''-Oxido dihydrogen<br />Light Water
| Section1 = {{Chembox Identifiers
|   CASNo = 7732-18-5
|   RTECS = ZC0110000
  }}
| Section2 = {{Chembox Properties
|   Formula = H<sub>2</sub>O or HOH
|   MolarMass = 18.01524&nbsp;g/mol
|   Appearance = transparent, almost<br />colorless liquid with<br />a slight hint of blue<ref>{{cite journal|last=Braun|first=Charles L.|coauthors=Sergei N. Smirnov|title=Why is water blue?|journal=J. Chem. Educ.|volume=70|issue=8|pages=612|date=1993|url=http://www.dartmouth.edu/~etrnsfer/water.htm|format=HTML}}</ref>
|   Density = 1000&nbsp;kg·m<sup>−3</sup>, liquid (4&nbsp;°C)<br /> 917&nbsp;kg·m<sup>−3</sup>, solid
|   MeltingPt = 0&nbsp;[[Celsius|°C]], 32&nbsp;[[Fahrenheit|°F]] (273.15&nbsp;[[kelvin|K]])<ref name="VSMOW">[[Vienna Standard Mean Ocean Water]] (VSMOW), used for calibration, melts at 273.1500089(10)&nbsp;K (0.000089(10)&nbsp;°C, and boils at 373.1339&nbsp;K (99.9839&nbsp;°C)</ref>
|   BoilingPt = 100&nbsp;°C, 212&nbsp;°F (373.15&nbsp;K)<ref name="VSMOW"/>
|   pKa = 15.74<br />~35-36
|   pKb = 15.74
|   Viscosity = 0.001&nbsp;[[pascal second|Pa·s]] at 20&nbsp;°C
  }}
| Section3 = {{Chembox Structure
|   MolShape = non-linear bent
|   CrystalStruct = Hexagonal<br />''See [[ice]]''
|   Dipole = 1.85&nbsp;[[Debye|D]]
  }}
| Section7 = {{Chembox Hazards
|   MainHazards = see [[Dihydrogen monoxide hoax]]
|   NFPA-H = 0
|   NFPA-F = 0 
|   NFPA-R = 0
  }}
| Section8 = {{Chembox Related
|   Function = [[solvent]]s
|   OtherFunctn = [[acetone]]<br />[[methanol]]
|   OtherCpds = [[water vapor]]<br />[[ice]]<br />[[heavy water]]
  }}
}}


'''Water''' ([[hydrogen|H]]<sub>2</sub>[[oxygen|O]], H[[Hydroxide|OH]]) is the most abundant [[molecule]] on [[Earth]]'s surface, composing of about 70% of the Earth's surface as liquid and solid state in addition to being found in the atmosphere as a vapor. It is in [[dynamic equilibrium]] between the [[liquid]] and [[vapor]] states at [[standard temperature and pressure]]. At [[room temperature]], it is a nearly [[colorless]], [[taste]]less, and [[odorless]] liquid, with a hint of blue.  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 clean, and may have some properties different from those in the laboratory.  However, there are many compounds that are essentially, if not completely, insoluble in water. Water is the only common, pure substance found naturally in all three common [[states of matter]]—for other substances, see [[water#Chemical and physical properties|Chemical properties]].
==Forms of water==
:''See the [[Water#Overview of types of water]]''
Water can take many forms.  The [[solid|solid state]] of water is commonly known as [[ice]] (while many other forms exist; see [[amorphous solid water]]); the [[gaseous state]] is known as [[water vapor]] (or [[steam]], though this is actually incorrect, since steam is just condensing ''liquid'' water droplets), and the common liquid [[phase (matter)|phase]] is generally taken as simply water.  Above a certain [[critical temperature]] and pressure (647 [[kelvin|K]] and 22.064 [[pascal (unit)|MPa]]), water molecules assume a ''supercritical'' condition, in which liquid-like clusters float within a vapor-like phase.

[[Heavy water]] is water in which the hydrogen is replaced by its heavier [[isotope]], [[deuterium]]. It is ''chemically'' almost identical to normal water.  Heavy water is used in the [[nuclear reactor|nuclear industry]] to slow down [[neutron]]s.

==Physics and chemistry of water==

Water is the [[chemical substance]] with [[chemical formula]] '''[[hydrogen|H]]<sub>2</sub>[[oxygen|O]]''': one [[molecule]] of water has two [[hydrogen]] [[atom]]s [[covalent]]ly [[chemical bond|bonded]] to a single [[oxygen]] atom. Water is a tasteless, odorless liquid at [[standard conditions|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.<ref>{{cite journal|last=Braun|first=Charles L.|coauthors=Sergei N. Smirnov|title=Why is water blue?|journal=J. Chem. Educ.|volume=70|issue=8|pages=612|date=1993|url=http://www.dartmouth.edu/~etrnsfer/water.htm|format=HTML}}</ref>
Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the [[Chalcogen|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 [[hydrogen]] to produce gases under standard conditions. The reason that oxygen hydride (water) forms a liquid is that it is more [[electronegative]] than all of these elements (other than 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]]. Water can be described as a polar liquid that dissociates disproportionately into the [[hydronium]] ion (H<sub>3</sub>O<sup>+</sup><sub>(aq)</sub>) and an associated [[hydroxide]] ion (OH<sup>−</sup><sub>(aq)</sub>).
Water is in [[dynamic equilibrium]] between the [[liquid]], [[gas]] and [[solid]] [[states of matter|states]] at [[standard temperature and pressure]], and is the only pure substance found naturally on Earth to be so. 

===Water, ice and vapor===
====Heat capacity and heat of vaporization====
{{main|Enthalpy of vaporization}}
Water has the second highest [[specific heat capacity]] of any known chemical compound, after [[ammonia]], as well as a high [[heat of vaporization]] (40.65 kJ mol<sup>−1</sup>), both of which are a result of the extensive [[hydrogen bond]]ing between its molecules. These two unusual properties allow water to moderate Earth's [[climate]] by buffering large fluctuations in temperature.

====Density of water and ice====
The solid form of most substances is more [[density|dense]] than the liquid [[phase (matter)|phase]]; thus, a block of pure solid ''substance'' will sink in a tub of pure liquid ''substance''.  But, by contrast, a block of common [[ice]] will float in a tub of water because solid water is ''less'' dense than liquid water.  This is an extremely important characteristic property of water.  At [[room temperature]], liquid water becomes denser with lowering temperature, just like other substances. But at 4 °C, just above freezing, water reaches its [[maximum density of water|maximum density]], and as water cools further toward its freezing point, the liquid water, under standard conditions, expands to become ''less'' dense.  The physical reason for this is related to the [[crystal structure]] of ordinary [[ice]], known as [[hexagonal (crystal system)|hexagonal]] [[ice Ih]]. Water, [[lead]], [[uranium]], [[neon]] and [[silicon]] are some of the few materials which expand when they freeze; most other materials contract. It should be noted however, that not all forms of ice are less dense than liquid water.  For example [[high density amorphous ice|HDA]] and [[very high density amorphous ice|VHDA]] are both more dense than liquid phase pure water.  Thus, the reason that the common form of ice is less dense than water is a bit non-intuitive and relies heavily on the unusual properties inherent to the [[hydrogen bond]].

Generally, water expands when it freezes because of its [[molecular structure]], in tandem with the unusual [[Elasticity (physics)|elasticity]] of the hydrogen bond and the particular lowest energy hexagonal [[crystal]] conformation that it adopts under standard conditions.  That is, when water cools, it tries to stack in a [[crystalline lattice]] configuration that stretches the [[rotational]] and [[oscillation|vibration]]al components of the bond, so that the effect is that each molecule of water is pushed further from each of its neighboring molecules.  This effectively reduces the density ''ρ'' of water when ice is formed under standard conditions.

The importance of this property cannot be overemphasized for its role on the [[ecosystem]] of Earth.  For example, ''if'' water were more dense when frozen, lakes and oceans in a polar environment would eventually freeze solid (from top to bottom). This would happen because frozen ice would settle on the lake and riverbeds, and the necessary warming phenomenon (see below) could not occur in summer, as the warm surface layer would be less dense than the solid frozen layer below.  It is a significant feature of nature that this does not occur naturally in the environment.

Nevertheless, the unusual expansion of freezing water (in ordinary ''natural'' settings in relevant biological systems), due to the [[hydrogen bond]], from 4 °C above freezing to the freezing point offers an important advantage for freshwater life in winter. Water chilled at the surface increases in density and sinks, forming [[convection]] [[current (fluid)|currents]] that cool the whole water body, but when the temperature of the lake water reaches 4 °C, water on the surface decreases in density as it chills further and remains as a surface layer which eventually [[freezing|freezes]] and forms ice. Since downward convection of colder water is blocked by the density change, any large body of fresh water frozen in winter will have the coldest water near the surface, away from the [[riverbed]] or lakebed. This accounts for various little known phenomena of ice characteristics as they relate to ice in lakes and "ice falling out of lakes" as described by early 20th century scientist Horatio D. Craft.

The following table gives the density of water in grams per cubic centimeter at various temperatures in degrees Celsius:<ref>Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Edn.). Boca Raton (FL):CRC Press.</ref>

{| class="wikitable"
!Temp (°C)!!Density (g/cm³)
|-
|30||0.9956502
|-
|25||0.9970479
|-
|22||0.9977735
|-
|20||0.9982071
|-
|15||0.9991026
|-
|10||0.9997026
|-
|4||0.9999720
|-
|0||0.9998395
|-
|−10||0.998117
|-
|−20||0.993547
|-
|−30||0.983854
|}
The values below 0 °C refer to [[supercooling|supercooled]] water.

<small>
;Freezing point
A simple but environmentally important and unusual property of water is that its usual solid form, [[ice]], floats on its liquid form. This solid state is not as dense as liquid water because of the geometry of the hydrogen bonds which are formed only at lower temperatures. For almost all other substances the solid form has a greater [[density]] than the liquid form. Fresh water at standard atmospheric pressure is most dense at 3.98&nbsp;°C, and will sink by [[convection]] as it cools to that temperature, and if it becomes colder it will rise instead. This reversal will cause deep water to remain warmer than shallower freezing water, so that ice in a body of water will form first at the surface and progress downward, while the majority of the water underneath will hold a constant 4&nbsp;°C. This effectively insulates a lake floor from the cold.
The water will freeze at 0&nbsp;°C (32&nbsp;°F, 273&nbsp;K), however, it can be [[supercooled]] in a fluid state down to its [[nucleation|crystal homogeneous nucleation]] at almost 231&nbsp;K (−42&nbsp;°C)<ref>P. G. Debenedetti, P. G., and Stanley, H. E.; "Supercooled and Glassy Water", Physics Today 56 (6), p. 40–46 (2003).</ref>.
Ice also has a number of more exotic phases not commonly seen (go to the full article on [[Ice]]).
</small>

====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 more dense 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 the forming ice causes it to sink towards the bottom.

====Miscibility and condensation====
{{main|Humidity}}
Water is [[miscible]] with many liquids, for example [[ethanol]] in all proportions, forming a single homogeneous liquid.  On the other hand water and most [[oil]]s are ''immiscible'' usually forming layers according to increasing density from the top.

[[Image:Relative Humidity.png|thumb|200px|right|Red line shows saturation]]
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]]''<ref>
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]]).
</ref> 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 ''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.<ref>
''[[Adiabatic process#Adiabatic heating and cooling|Adiabatic cooling]]'' resulting from the [[ideal gas law]].
</ref>

====Vapor Pressures of Water====
{| class="wikitable"
!Temperature (°C)
!Pressure (torr)
|-
|0
|4.58
|-
|5
|6.54
|-
|10
|9.21
|-
|12
|10.52
|-
|14
|11.99
|-
|16
|13.63
|-
|17
|14.53
|-
|18
|15.48
|-
|19
|16.48
|-
|20
|17.54
|-
|21
|18.65
|-
|22
|19.83
|-
|23
|21.07
|-
|24
|22.38
|-
|25
|23.76
|}
<ref>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.</ref>

====Compressibility====
The [[compressibility]] of water is a function of pressure and temperature. At 0&nbsp;°C in the limit of zero pressure the compressibility is 5.1&times;10<sup>-5</sup>&nbsp;bar<sup>−1</sup>.<ref>{{cite journal |author=Fine, R.A. and Millero, F.J. |year=1973 |title=Compressibility of water as a function of temperature and pressure |volume=59 |issue=10 |page=5529 |journal=Journal of Chemical Physics |doi=10.1063/1.1679903 }}</ref>  In the zero pressure limit the compressibility reaches a minimum of 4.4&times;10<sup>-5</sup>&nbsp;bar<sup>−1</sup> around 45&nbsp;°C before increasing again with increasing temperature.  As the pressure is increased the compressibility decreases, being 3.9&times;10<sup>-5</sup>&nbsp;bar<sup>−1</sup> at 0&nbsp;°C and 1000&nbsp;[[Bar (unit)|bar]].
The [[bulk modulus]] of water is 2.2&times;10<sup>9</sup>&nbsp;[[pascal (unit)|Pa]].<ref name = nave>{{cite web | title = Bulk Elastic Properties | author = R. Nave | work = HyperPhysics | publisher = [[Georgia State University]] | url = http://hyperphysics.phy-astr.gsu.edu/hbase/hph.html | accessdate = 2007-10-26}}</ref> 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 [[ocean]]s at 4000 [[metre|m]] depth, where pressures are 4×10<sup>7</sup>&nbsp;Pa, there is only a 1.8% decrease in volume.<ref name = nave/>

====Triple point====
{{main|Triple point}}
{| class="wikitable" style="float:right;"
|+The various triple points of water<ref name=Schleuter>{{cite paper|title=Impact of High Pressure &mdash; Low Temperature Processes on Cellular Materials Related to Foods|author=Oliver Schlüter|publisher=Technischen Universität Berlin|url=http://edocs.tu-berlin.de./diss/2003/schlueter_oliver.pdf|format=PDF|date=[[2003-07-28]]}}</ref>
!Phases in stable equilibrium
!Pressure
!Temperature
|-
|liquid water, [[ice I]], and water vapour
|611.73 Pa
|273.16 K
|-
|liquid water, [[ice Ih]], and [[ice III]]
|209.9 MPa
|251 K (-22 °C)
|-
|liquid water, ice Ih, and gaseous water
|612 Pa
|0.01 °C
|-
|liquid water, ice III, and [[ice V]]
|350.1 MPa
| -17.0 °C
|-
|liquid water, ice V, and [[ice VI]]
|632.4 MPa
|0.16 °C
|-
|ice Ih, [[Ice II]], and ice III
|213 MPa
| -35 °C
|-
|ice II, ice III, and ice V
|344 MPa
| -24 °C
|-
|ice II, ice V, and ice VI
|626 MPa
| -70 °C
|}

The [[temperature]] and [[pressure]] at which solid, liquid, and [[water vapor|gaseous water]] coexist in equilibrium is called the [[Triple point#Triple point of water|triple point of water]].  This point is used to define the units of temperature (the [[kelvin]] and, indirectly, the degree [[Celsius]] and even the degree [[Fahrenheit]]).  The triple point is at a temperature of 273.16 K (0.01 °C) by convention, and at a pressure of 611.73 [[pascal (unit)|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 is defined by the height at which the atmospheric pressure corresponds to the triple point of water.

[[Image:Water phase diagram.svg|right|400px|thumb|water [[phase diagram]]</br>
Y-axis = Pressure in Pascal (10<sup>n</sup>),</br>
X-axis = Temperature in Kelvin.</br>
S = Solid</br>
L = Liquid</br>
V = Vapour</br>
CP = Critical Point</br>
TP = [[Triple point#Triple point of water|Triple point of water]]]]

<small>
The [[triple point]] of water (the single combination of pressure and temperature at which pure liquid water, ice, and water vapor can coexist in a stable equilibrium) is used to define the [[kelvin]], the SI unit of thermodynamic temperature. As a consequence, water's triple point temperature is a prescribed value rather than a measured quantity: 273.16&nbsp;kelvins (0.01&nbsp;°C) and a pressure of 611.73&nbsp; pascals (approximately 0.0060373&nbsp;[[atmosphere (unit)|atm]]).
This is approximately the combination that exists with 100% relative humidity at sea level and the freezing 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.<ref>{{cite paper|title=The States Of Aggregation|date=1925|author=Gustav Heinrich Johann Apollon Tammann|publisher=Constable And Company Limited}}</ref><ref name=Schleuter /><ref>{{cite book|title=A System of Physical Chemistry|author=William Cudmore McCullagh Lewis and James Rice|date=1922|publisher=Longmans, Green and co.}}</ref>
</small>

====Mpemba effect====
The [[Mpemba effect]] is the surprising phenomenon whereby hot water can, under certain conditions, freeze sooner than cold water, even though it must pass the lower temperature on the way to freezing.  However, this can be explained with [[evaporation]], [[convection]], [[supercooling]], and the [[thermal insulation|insulating]] effect of [[frost]].

====Hot ice====
Hot ice is the name given to another surprising phenomenon in which water at room temperature can be turned into ice ''that remains at room temperature'' by supplying an electric field on the order of 10<sup>6</sup> volts per meter.<ref>{{cite journal | author=Choi, Eun-Mi; Yoon, Young-Hwan; Lee, Sangyoub; Kang, Heon | title = Freezing Transition of Interfacial Water at Room Temperature under Electric Fields | journal = [[Physical Review Letters]] | volume = 95 | issue = 8 | pages = 085701 | doi = 10.1103/PhysRevLett.95.085701}}</ref>

The effect of such electric fields has been suggested as an explanation of cloud formation. The first time cloud ice forms around a clay particle, it requires a temperature of −10 °C, but subsequent freezing around the same clay particle requires a temperature of just −5 °C, suggesting some kind of structural change.<ref>{{cite journal |title=Aircraft observations of the influence of electric fields on the aggregation of ice crystals |author=Connolly PJ, Saunders CPR, Gallagher MW, Bower KN, Flynn MJ, Choularton TW, Whiteway J, Lawson RP |journal=Quarterly Journal of the Royal Meteorological Society, Part B |year=2005 |month=April |volume=131 |issue=608 |pages=1695–1712}}</ref>

===Surface tension===
Water drops are stable, due to the high [[surface tension]] of water, 72.8 mN/m, the highest of the non-metallic liquids. This can be seen when small quantities of water are put on a surface such as glass: the water stays together as drops. This property is important for life. For example, when water is carried through [[xylem]] up stems in plants the strong intermolecular attractions hold the water column together. Strong cohesive properties hold the water column together, and strong adhesive properties stick the water to the xylem, and prevent tension rupture caused by [[transpiration pull]]. Other liquids with lower surface tension would have a higher tendency to "rip", forming vacuum or air pockets and rendering the xylem water transport inoperative.

===Electrical properties===
Pure water containing no ions is an excellent [[Electrical insulation|insulator]], however, not even "deionized" water, is completely free of ions. Water undergoes [[self-ionization of water|auto-ionisation]] at any temperature above [[absolute zero]].  Further, 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 [[ion]]s in aqueous solution by which an electric current can flow.

Water can be split into its constituent elements, hydrogen and oxygen, by passing a current through it. This process is called [[electrolysis]]. Water molecules naturally dissociate into H<sup>+</sup> and OH<sup>−</sup> ions, which are pulled toward the [[cathode]] and [[anode]], respectively. At the cathode, two H<sup>+</sup> ions pick up electrons and form H<sub>2</sub> gas. At the anode, four OH<sup>−</sup> ions combine and release O<sub>2</sub> gas, molecular water, and four electrons. The gases produced bubble to the surface, where they can be collected. It is known that the theoretical maximum electrical resistivity for water is approximately 182 [[kilohm|k&Omega;]]·m²/m (or 18.2 MΩ·cm²/cm) at 25 °C. This figure agrees well with what is typically seen on [[reverse osmosis]], ultrafiltered and deionized [[ultrapure water]] systems used for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding that of even 100 parts per trillion (ppt) in ultrapure water will begin to noticeably lower its resistivity level by up to several kilohm-square meters per meter (a change of several hundred [[siemens (unit)|nanosiemens]] per meter of conductance).

<small>
;Electrical conductivity
Pure water has a ''low'' [[electrical conductivity]], but this increases significantly upon solvation of a small amount of ionic material water such as [[hydrogen chloride]]. Thus the risks of [[electric shock|electrocution]] are much greater in water with the usual impurities not found in pure water. Any electrical properties observable in water are from the [[ion]]s of mineral salts and [[carbon dioxide]] dissolved in it. [[self-ionization of water|Water does self-ionize]] where two water molecules become 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 [[electrical conductivity|conductivity]] of 0.055 [[Siemens (unit)|µS]]/[[Centimeter|cm]] at 25&nbsp;°C. Water can also be [[electrolysis|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 (and some other electrolytes), [[protons]] are the primary carriers (see [[proton conductor]]).
</small>

===Dipolar nature of water===
[[Image:3D model hydrogen bonds in water.jpg|right|thumb|model of [[hydrogen bond]]s between molecules of water]]
An important feature of water is its [[polar molecule|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. A molecule 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 is known as [[hydrogen bond]]ing, and explains many of the properties of water. Certain molecules, such as carbon dioxide, also have a difference in electronegativity between the atoms but the difference is that the shape of carbon dioxide is symmetrically aligned and so the opposing charges cancel one another out. This phenomenon of water can be seen if you hold an electrical source near a thin stream of water falling vertically, causing the stream to bend towards the electrical source.

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 point|melting]] and [[boiling point]] temperatures; more [[heat]] energy is required to break the hydrogen bonds between molecules. The similar compound hydrogen sulfide (H<sub>2</sub>S), 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.

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, will float in water. In other words, water expands as it freezes, whereas almost all other materials shrink on solidification.

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 &mdash;the pressure exerted by an ice skater on the ice would only reduce the melting point by approximately 0.09 °C (0.16 °F).

<small>
;Electronegative Polarity
Water has a partial negative charge (σ-) near the oxygen atom due to the unshared pairs of electrons, and partial positive charges (σ+) near the hydrogen atoms. In water, this happens because the oxygen atom is more [[electronegative]] than the hydrogen atoms — that is, it has a stronger "[[electrostatic force|pulling power]]" on the molecule's [[electron]]s, drawing them closer (along with their negative charge) and making the area around the oxygen atom more negative than the area around both of the hydrogen atoms. 

====Adhesion====
[[Image:Water drops on spider web.jpg|thumb|right|[[Dew]] drops adhering to a [[spider web]]]]
Water sticks to itself ([[cohesion (chemistry)|cohesion]]) because it is [[polar molecule|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 between glass and water molecules (adhesive forces) are stronger than the cohesive forces.
In biological cells and [[organelle]]s, 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&mdash;to remove the strongly held layers of water of hydration&mdash;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]].<ref> [http://www.biophysics.org/education/parsegian.pdf Physical Forces Organizing Biomolecules (PDF)]</ref> They are particularly important when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.

====Surface tension====
{{main|Surface tension}}
[[Image:Dscn3156-daisy-water 1200x900.jpg|thumb|right|This [[daisy]] is under the water level, which has risen gently and smoothly. Surface tension prevents the water from submerging the flower.]]

Water has a high [[surface tension]] caused by the strong cohesion between water molecules. This can be seen when small quantities of water are put onto a non-soluble surface such as [[polythene]]; 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 wave]]s which are the surface ripples that form from around the impact of drops on water surfaces, and some times occur with strong subsurface currents flow to the water surface. The apparent elasticity caused by surface tension drives the waves.

====Capillary action====
{{main|Capillary action}}
[[Capillary action]] refers to the process of water moving up a narrow tube against the force of [[gravity]]. It occurs because water adheres to the sides of the tube, and then surface tension tends to straighten the surface making the surface rise, and more water is pulled up through cohesion. The process is repeated as the water flows up the tube until there is enough water that gravity can counteract the adhesive force.
</small>

===Water as a solvent===
Water is also a good [[solvent]] due to its [[polar molecule|polarity]]. When an ionic or polar compound enters water, it is surrounded by water molecules ([[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 [[acid]]s, [[alcohol]]s, and [[salt]]s 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 force|van der Waals interactions]]  with nonpolar molecules.

An example of an ionic solute is [[sodium chloride|table salt]]; the sodium chloride, NaCl, separates into Na<sup>+</sup> [[cation]]s and Cl<sup>-</sup> [[anion]]s, each being surrounded by water molecules. The ions are then easily transported away from their [[crystal lattice|crystalline lattice]] into solution. An example of a nonionic solute is [[sucrose|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.

<small>
;Solvation
[[Image:Havasu Falls 2 md.jpg|thumb|200px|left|High concentrations of dissolved [[Lime (mineral)|lime]] make the water of [[Havasu Falls]] appear turquoise.]]
Water is a very strong [[solvent]], referred to as ''the universal solvent'', dissolving many types of substances. Substances that will mix well and dissolve in water (e.g. [[Salt (chemistry)|salt]]s) are known as "[[hydrophilic]]" (water-loving) substances, while those that do not mix well with water (e.g. [[lipids|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 [[intermolecular force#Dipole-dipole interactions|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 "[[precipitation (chemistry)|pushed out]]" from the water, and do not dissolve. Contrary to the common misconception, water and hydrophobic substances does not "repel", and the hydration of a hydrophobic surface is energetically, but not entropically, favorable.
</small>

===Amphoteric nature of water===
Chemically, water is [[amphoteric]] — i.e., it is able to act as either an [[acid]] or a [[base (chemistry)|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<sup>−</sup>) is equal to that of the [[hydronium]] (H<sub>3</sub>O<sup>+</sup>) or [[hydrogen]] (H<sup>+</sup>) ions. If the [[chemical equilibrium|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 [[Brønsted-Lowry]] system, an acid is defined as a species which donates a proton (an H<sup>+</sup> 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 instance, it receives an H<sup>+</sup> ion from HCl in the equilibrium:

:HCl + H<sub>2</sub>O {{unicode|⇌}} H<sub>3</sub>O<sup>+</sup> + Cl<sup>−</sup>

Here water is acting as a base, by receiving an H<sup>+</sup> ion.

In the reaction with [[ammonia]], NH<sub>3</sub>, water donates an H<sup>+</sup> ion, and is thus acting as an acid:

:NH<sub>3</sub> + H<sub>2</sub>O {{unicode|⇌}} NH<sub>4</sub><sup>+</sup> + OH<sup>−</sup>

===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 CO<sub>2</sub> 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.

===Hydrogen bonding in water===
A water molecule can form a maximum of four [[hydrogen bond]]s because it can accept two and donate two hydrogens. Other molecules like [[hydrogen fluoride]], [[ammonia]], [[methanol]] form hydrogen bonds but they do not show anomalous behaviour of [[thermodynamics|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 [[liquid]]s. 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 temperature regime that suits diverse [[biological process]]es, 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.

===Quantum properties of molecular water===
Although the molecular formula of water is generally considered to be a stable result in molecular thermodynamics, recent work started in 1995 has shown that at certain scales, water may act more like H<sub>3/2</sub>O than H<sub>2</sub>O at the quantum level.<ref>{{cite news | publisher = Physics News Update |date=31 Jul 03 | author = Phil Schewe, James Riordon, and Ben Stein | title = A Water Molecule's Chemical Formula is Really Not H<sub>2</sub>O | url = http://www.aip.org/enews/physnews/2003/split/648-1.html}}</ref> This result could have significant ramifications at the level of, for example, the [[hydrogen bond]] in [[biology|biological]], [[chemistry|chemical]] and [[physics|physical]] systems.  The experiment shows that when [[neutrons]] and [[electrons]] collide with water, they scatter in a way that indicates that they only are affected by a ratio of 1.5:1 of [[hydrogen]] to [[oxygen]] respectively.  However, the time-scale of this response is only seen at the level of attoseconds (10<sup>-18</sup> seconds), and so is only relevant in highly resolved [[kinetic]] and [[dynamics (mechanics)|dynamical]] systems.<ref>{{cite journal | author = C. A. Chatzidimitriou-Dreismann, T. Abdul Redah, R. M. F. Streffer and J. Mayers | title = Anomalous Deep Inelastic Neutron Scattering from Liquid H<sub>2</sub>O-D<sub>2</sub>O: Evidence of Nuclear Quantum Entanglement | year = 1997 | journal = [[Physical Review Letters]] | volume = 79 | issue = 15 | pages = 2839 | doi = 10.1103/PhysRevLett.79.2839}}</ref><ref>{{cite journal | author = C. A. Chatzidimitriou-Dreismann, M. Vos, C. Kleiner and T. Abdul-Redah | title = Comparison of Electron and Neutron Compton Scattering from Entangled Protons in a Solid Polymer | year = 2003 | journal = [[Physical Review Letters]] | volume = 91 | issue = 5 | pages = 057403-4 | doi = 10.1103/PhysRevLett.91.057403}}</ref>

=== Heavy Water and [[isotopologues]] of water ===
Hydrogen has three isotopes. The most common, making up more than 95% of water, has 1 proton and 0 neutrons. A second isotope, [[deuterium]] (short form "D"), has 1 proton and 1 neutron. Deuterium, {{chem|D|2|O}}, is also known as [[heavy water]] and is used in [[nuclear reactor]]s as a [[neutron moderator]].  The third isotope, [[tritium]], has 1 proton and 2 neutrons, and is [[radioactive]], with a [[half-life]] of 12.32 years. {{chem|T|2|O}} exists in nature only in tiny quantities, being produced primarily via cosmic ray-driven nuclear reactions in the atmosphere. {{chem|D|2|O}} is stable, but differs from {{chem|H|2|O}} in in that it is more dense - hence, "heavy water" - and in that several other physical properties are slightly different from those of common, Hydrogen-1 containing "light water". {{chem|D|2|O}} occurs naturally in ordinary water in very low concentrations. Consumption of pure isolated {{chem|D|2|O}} 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.

===Transparency===
Water's transparency is also an important property of the liquid. If water were not transparent, sunlight, essential to aquatic plants, would not reach into seas and oceans.

==History==

The properties of water have historically been used to define various [[Temperature conversion|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.<ref>http://home.comcast.net/~igpl/Temperature.html</ref>

The first scientific decomposition of water into hydrogen and oxygen, by [[electrolysis]], was done in 1800 by [[William Nicholson (chemist)|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 [[polymer]]ized form of water that was the subject of much scientific controversy during the late 1960s.  The consensus now is that it does not exist.

[[Pseudoscience]] concept is [[water memory]].

==Systematic naming==
The accepted [[IUPAC nomenclature of inorganic chemistry|IUPAC]] name of water is simply "water", 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 oxygen-based [[substituent group]]s,<ref>Leigh, G. J. ''et al.'' 1998. [http://www.iupac.org/publications/books/principles/principles_of_nomenclature.pdf ''Principles of chemical nomenclature: a guide to IUPAC recommendations''], p. 99. Blackwell Science Ltd, UK. ISBN 0-86542-685-6</ref> 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]].

===Systematic nomenclature and humor===
{{main|Dihydrogen monoxide hoax}}

'''Dihydrogen monoxide''' or '''DHMO''' is an overly pedantic systematic covalent name of water. This term has been used in [[parody|parodies]] of chemical research that call for this "lethal chemical" to be banned. In reality, a more realistic systematic name would be '''hydrogen oxide''', since the "di-" and "mon-" prefixes are superfluous. [[Hydrogen sulfide]], H<sub>2</sub>S, is never referred to as "dihydrogen monosulfide", and [[hydrogen peroxide]], H<sub>2</sub>O<sub>2</sub>, is never called "dihydrogen dioxide".

Some overzealous [[material safety data sheet]]s for water list the following:  Caution: May cause drowning!{{Fact|date=October 2007}}

Other systematic names for water include '''hydroxic acid''' or '''hydroxylic acid'''. Likewise, the systematic alkali name of water is '''hydrogen hydroxide'''&mdash;both acid and alkali names exist for water because it is able to react both as an acid or an alkali, depending on the strength of the acid or alkali it is reacted with ([[amphoterism|amphoteric]]).  None of these names are used widely outside of DHMO sites.

==See also==
*[[Double distilled water]]
*[[Hydrodynamics]]
*[[Viscosity#Viscosity of water|Viscosity of Water]]
*[[Vienna Standard Mean Ocean Water]]
*[[Water absorption]] of [[electromagnetic radiation]]
*[[Water dimer]]
*[[Water (data page)]]
</div>

==References==
<div class="references-small">
<references/>
</div>

==External links==
*[http://www.iapws.org/relguide/IF97.pdf Release on the IAPWS Industrial Formulation 1997 for the Thermodynamic Properties of Water and Steam] (fast computation speed)
*[http://www.iapws.org/relguide/IAPWS95.pdf Release on the IAPWS Formulation 1995 for the Thermodynamic Properties of Ordinary Water Substance for General and Scientific Use] (simpler formulation)
*[http://www.dhmo.org/ A spoof site on the "dangers" of dihydrogen monoxide]
*[http://www.siwi.org/ Stockholm International Water Institute] (SIWI)
*[http://www.btinternet.com/~martin.chaplin/explan.html Explanation of the anomalous properties of water]
*[http://www.compchemwiki.org/index.php?title=Water Computational Chemistry Wiki]
*[http://www.lsbu.ac.uk/water/phase.html Water phase diagrams]
*[http://www.lsbu.ac.uk/water/vibrat.html Water Absorption Spectrum]

[[Category:Forms of water]]
[[Category:Hydrides]]
[[Category:Hydrogen compounds]]
[[Category:Hydroxides]]
[[Category:Oxides]]
[[Category:Inorganic solvents]]
[[Category:Water chemistry| ]]
[[Category:Neutron moderators]]

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[[af:water (molekule)]]
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