water+molecule

=Please stop plagiarizing from Wikipedia!!!!!!!!!! Please fix what you have stolen.= =Water= The electronegativity difference tells us which molecule attracts more atoms. Because oxygen has a higher electronegativity than hydrogen, the oxygen atoms will attract more electrons. The electronegativity difference between atoms, changes the energy of that bond.
 * Water ** is a chemical substance with the chemical formula H 2 O. It is a bent structure and has an electronegativity difference of 1.4.

Water is a __special__ molecule because of it's **hydrogen bonds.** Hydrogen bonds are extremely strong and not easily broken. To be exact, O-H bonds have a strength of 467 KJ/mol. The higher the strength of a bond, the more energy the bond requires to be broken. Because the hydrogen bonds require such a high amount of energy to be broken, water has a very high specific heat, meaning, it takes a lot of energy to break hydrogen bonds. This is why water is useful for __heating and cooling__ other things; because, the water itself has a higher boiling and freezing point, so it will not break while trying to break the other bonds. Water hydrogen bonds are very unique in that, in its liquid state, the hydrogen bonds are very condensed, while in the solid state, hydrogen bonds create a more open space between bonds. This is why the solid state of water (ice) floats in the liquid state (water), unlike most other substances. Hydrogen bonds are not only limited to water, they also exist in bonds containing Nitrogen and Fluorine.

Water is a polar molecule where the positive Hydrogen and the negative Oxygen, attract, making it a polar molecule. That __charge__ separation is due to the electronegativity difference. Because water is a polar molecule, following the like dissolves like rule, dissolves other polar molecules.

Water is used for four major categories: thermal electric power, irrigation, domestic and industrial.

A water molecule contains one oxygen and two hydrogen atoms connected by covalent bonds. Water often co-exists on Earth with its solid, (ice), or in a gas state of water vapor or steam. Under the nomenclature used to name chemical compounds // Dihydrogen monoxide // is the scientific name for water, though it is almost never used. Water boils at 100 __degrees__ Celsius and freezes at 0 degrees Celsius.

Water also has special properties that allow its liquid form to be more dense than its solid form. Click here to find out more.




 * [|IUPAC name] [|[hide]] Water

Oxidane ||
 * Other names [|[hide]] Hydrogen oxide

[|Dihydrogen monoxide]

Hydrogen monoxide

Dihydrogen oxide

Hydrogen hydroxide

Hydrol [|[1]] ||  [|[show]] ||  [|[show]] || 917 kg/m3, solid || ~35–36 || [|Water intoxication] ||  000 || [|Hydrogen selenide] [|Hydrogen telluride] [|Hydrogen polonide] [|Hydrogen peroxide] || [|methanol] || [|ice] [|heavy water] ||
 * ~ Identifiers ||
 * [|CAS number] || [|7732-18-5]  [[image:http://upload.wikimedia.org/wikipedia/en/thumb/f/fb/Yes_check.svg/7px-Yes_check.svg.png width="7" height="7" caption="Yes"]] ||
 * [|PubChem] || [|962]  ||
 * [|ChemSpider] || [|937]  [[image:http://upload.wikimedia.org/wikipedia/en/thumb/f/fb/Yes_check.svg/7px-Yes_check.svg.png width="7" height="7" caption="Yes"]] ||
 * [|UNII] || [|059QF0KO0R]  [[image:http://upload.wikimedia.org/wikipedia/en/thumb/f/fb/Yes_check.svg/7px-Yes_check.svg.png width="7" height="7" caption="Yes"]] ||
 * [|ChEBI] || [|CHEBI:15377]  [[image:http://upload.wikimedia.org/wikipedia/en/thumb/f/fb/Yes_check.svg/7px-Yes_check.svg.png width="7" height="7" caption="Yes"]] ||
 * [|ChEMBL] || [|CHEMBL1098659]  [[image:http://upload.wikimedia.org/wikipedia/en/thumb/f/fb/Yes_check.svg/7px-Yes_check.svg.png width="7" height="7" caption="Yes"]] ||
 * [|RTECS number] || ZC0110000 ||
 * [|Jmol] -3D images || [|Image 1] ||
 * [|SMILES]
 * [|InChI]
 * ~ Properties ||
 * [|Molecular formula] || H 2 O ||
 * [|Molar mass] || 18.01528(33) g/mol ||
 * Appearance || white solid or almost colorless, transparent, with a slight hint of blue, crystalline solid or liquid [|[2]] ||
 * [|Density] || 1000 kg/m3, liquid (4 °C) (62.4 lb/cu. ft)
 * [|Melting point] || 0 [|°C], 32 ° [|F] , (273.15 [|K] ) [|[3]] ||
 * [|Boiling point] || 99.98 °C, 211.97 °F (373.13 K) [|[3]] ||
 * [|Acidity] (p//K//a) || 15.74
 * [|Basicity] (p//K//b) || 15.74 ||
 * [|Refractive index] (//n//D) || 1.3330 ||
 * [|Viscosity] || 0.001 [|Pa s] at 20 °C ||
 * ~ Structure ||
 * [|Crystal structure] || [|Hexagonal] ||
 * [|Molecular shape] || [|Bent] ||
 * [|Dipole moment] || 1.85 [|D] ||
 * ~ Hazards ||
 * Main [|hazards] || Drowning (see also [|Dihydrogen monoxide hoax] )
 * [|NFPA 704] || [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/6/6f/NFPA_704.svg/75px-NFPA_704.svg.png width="75" height="75" caption="NFPA 704.svg" link="http://en.wikipedia.org/wiki/File:NFPA_704.svg"]]
 * ~ Related compounds ||
 * Other [|cations] || [|Hydrogen sulfide]
 * Related [|solvents] || [|acetone]
 * Related compounds || [|water vapor]

Forms of water
Like many substances, [|water] can take numerous forms that are broadly categorized by [|phase of matter]. The [|liquid phase] is the most common among water's phases (within the Earth's atmosphere and surface) and is the form that is generally denoted by the word "water." The [|solid phase] of water is known as [|ice] and commonly takes the structure of hard, amalgamated [|crystals], such as [|ice cubes] , or loosely accumulated [|granular] crystals, like [|snow]. For a list of the many different crystalline and [|amorphous] forms of solid H 2 O, see the article [|ice]. The [|gaseous phase] of water is known as [|water vapor] (or [|steam] ), and is characterized by water assuming the configuration of a transparent [|cloud]. (Note that the visible steam and clouds are, in fact, water in the liquid form as minute droplets suspended in the air.) The fourth state of water, that of a [|supercritical fluid], is much less common than the other three and only rarely occurs in nature, in extremely uninhabitable conditions. When water achieves a specific [|critical temperature] and a specific [|critical pressure] (647  [|K] and 22.064  [|MPa] ), liquid and gas phase merge to one homogeneous fluid phase, with properties of both gas and liquid. One example of naturally occurring supercritical water is found in the hottest parts of deep water [|hydrothermal vents], in which water is heated to the critical temperature by scalding [|volcanic] [|plumes] and achieves the critical pressure because of the crushing __weight__ of the ocean at the extreme depths at which the vents are located. Additionally, anywhere there is volcanic activity below a depth of 2.25 km (1.40 mi) can be expected to have water in the supercritical phase. [|[6]] [|Vienna Standard Mean Ocean Water] is the current international standard for water [|isotopes]. Naturally occurring water is almost completely composed of the neutron-less hydrogen isotope [|protium]. Only 155 [|ppm] include [|deuterium] ( 2H or D), a hydrogen isotope with one neutron, and fewer than 20 parts per [|quintillion] include [|tritium] ( 3H or T), which has two. ** [|Heavy water] ** is water with a higher-than-average deuterium content, up to 100%. Chemically, it is similar but not identical to normal water. This is because the nucleus of deuterium is twice as heavy as protium, and this causes noticeable differences in bonding energies. Because water molecules exchange hydrogen atoms with one another, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure dideuterium monoxide (D 2 O). Humans are generally unaware of taste differences, [|[7]] but sometimes report a burning sensation [|[8]] or sweet flavor. [|[9]] Rats, however, are able to avoid heavy water by smell. [|[10]] Toxic to many animals, [|[10]] heavy water is used in the [|nuclear reactor] industry to [|moderate] (slow down) [|neutrons]. [|Light water reactors] are also common, where "light" simply designates normal water. ** [|Light water] ** more specifically refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard 155ppm level. Light water has been found to be beneficial for improving cancer survival rates in mice [|[11]] and humans undergoing chemotherapy

Physics and chemistry
See also: [|Water chemistry analysis] Water is the [|chemical substance] with [|chemical formula] H 2  O : one [|molecule] of water has two [|hydrogen] [|atoms] [|covalently] [|bonded] to a single [|oxygen] atom. [|[13]] 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. [|[2]] 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]. 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 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 timescales faster than 200 femtoseconds. [|[14]] However, this bond is sufficiently strong to create many of the peculiar properties of water, such as those that make it integral to life. Water can be described as a [|polar] liquid that slightly dissociates disproportionately into the [|hydronium] ion ( H 3 O +(aq)) and an associated [|hydroxide] ion (OH−(aq)). 2 H 2  O (l)  H  3  O+ (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

Heat capacity and heats of vaporization and fusion
Heat of vaporization||~ Temperature (°C) Main article: [|Enthalpy of vaporization]    Heat of vaporization of Water from melting to critical temperature Water has a very high [|specific heat capacity] – the second highest among all the heteroatomic species (after [|ammonia] ), as well as a high [|heat of vaporization] (40.65 kJ/mol or 2257 kJ/kg at the normal boiling point), 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. According to Josh Willis, of [|NASA] 's [|Jet Propulsion Laboratory], the oceans absorb one thousand times more heat than the atmosphere (air) and are holding 80 to 90% of [|global warming] heat. [|[16]] The specific [|enthalpy of fusion] of water is 333.55 kJ/kg at 0 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice of [|glaciers] and [|drift ice]. Before and since the advent of mechanical [|refrigeration], ice was and still is in common use for retarding food spoilage. Constant-pressure heat capacity||~ Temperature (°C)
 * ~ //H//v (kJ/mol) [|[15]] ||
 * 0 || 45.054 ||
 * 25 || 43.99 ||
 * 40 || 43.35 ||
 * 60 || 42.482 ||
 * 80 || 41.585 ||
 * 100 || 40.657 ||
 * 120 || 39.684 ||
 * 140 || 38.643 ||
 * 160 || 37.518 ||
 * 180 || 36.304 ||
 * 200 || 34.962 ||
 * 220 || 33.468 ||
 * 240 || 31.809 ||
 * 260 || 29.93 ||
 * 280 || 27.795 ||
 * 300 || 25.3 ||
 * 320 || 22.297 ||
 * 340 || 18.502 ||
 * 360 || 12.966 ||
 * 374 || 2.066 ||
 * ~ //C////p// (J/(g·K) at 100 kPa) [|[17]] ||
 * 0 || 4.2176 ||
 * 10 || 4.1921 ||
 * 20 || 4.1818 ||
 * 30 || 4.1784 ||
 * 40 || 4.1785 ||
 * 50 || 4.1806 ||
 * 60 || 4.1843 ||
 * 70 || 4.1895 ||
 * 80 || 4.1963 ||
 * 90 || 4.205 ||
 * 100 || 4.2159 ||

Note that the specific heat capacity of ice at −10 °C is [|about 2.05] J/(g·K) and that the heat capacity of steam at 100 °C is [|about 2.080] J/(g·K).

Density of water and ice
   Density of ice and water as a function of temperature Density of liquid water||~ Temp (°C) The density of water is approximately one gram per cubic centimeter. More precisely, it is dependent on its temperature, but the relation is not linear and is [|unimodal] rather than [|monotonic] (see right-hand table). When cooled from [|room temperature] liquid water becomes increasingly dense, just like other substances. But at approximately 4 °C (39 °F), pure water reaches its [|maximum density]. As it is cooled further, it expands to become less dense. This unusual negative thermal expansion is attributed to strong, orientation-dependent, intermolecular interactions and is also observed in molten [|silica]. [|[20]] The solid form of most substances is [|denser] than the liquid [|phase] ; thus, a block of most solids will sink in the liquid. However, a block of ice floats in liquid water because ice is //less// dense. Upon freezing, the density of water decreases by about 9%. [|[21]] The reason for this is the 'cooling' of intermolecular vibrations allowing the molecules to form steady hydrogen bonds with their neighbors and thereby gradually locking into positions reminiscent of the [|hexagonal] packing achieved upon freezing to [|ice Ih]. Whereas the hydrogen bonds are shorter in the crystal than in the liquid, this locking effect reduces the average coordination number of molecules as the liquid approaches nucleation. Other substances that expand on freezing are [|silicon], [|gallium] , [|germanium] , [|antimony] , [|bismuth] , [|plutonium] and other compounds that form spacious crystal lattices with tetrahedral coordination. Only ordinary hexagonal ice is less dense than the liquid. Under increasing pressure, ice undergoes a number of transitions to other [|allotropic forms] with higher density than liquid water, such as [|high density amorphous ice] (HDA) and [|very high density amorphous ice] (VHDA).    Temperature distribution in a lake in summer and winter Water also expands significantly as the temperature increases. Its density decreases by 4% from its highest value when approaching its boiling point. The melting point of ice is 0 °C (32 °F, 273.15 K) at standard pressure, however, pure liquid water can be [|supercooled] well below that temperature without freezing if the liquid is not mechanically disturbed. It can remain in a fluid state down to its homogeneous [|nucleation] point of approximately 231 K (−42 °C). [|[22]] The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, but as ice transforms into its [|allotropes] (see [|crystalline states of ice] ) above 209.9 MPa (2,072 atm), the melting point increases markedly [|with pressure], i.e., reaching 355 K (82 °C) at 2.216 GPa (21,870 atm) (triple point of [|Ice VII][|[23]] ). A significant increase of pressure is required to lower the melting point of ordinary ice—the pressure exerted by an ice skater on the ice only reduces the melting point by approximately 0.09 °C (0.16 °F).[// [|citation needed] //] These properties of water have important consequences in its role in the [|ecosystem] of Earth. Water at 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 [|[24]] (good insulators) it is unlikely that sufficiently deep lakes will freeze completely, unless stirred by strong currents that mix cooler and warmer water and accelerate the cooling. In warming weather, chunks of ice float, rather than sink to the bottom where they might melt extremely slowly. These phenomena thus may help to preserve aquatic life.
 * ~ Density (kg/m3) [|[18]][|[19]] ||
 * +100 || 958.4 ||
 * +80 || 971.8 ||
 * +60 || 983.2 ||
 * +40 || 992.2 ||
 * +30 || 995.6502 ||
 * +25 || 997.0479 ||
 * +22 || 997.7735 ||
 * +20 || 998.2071 ||
 * +15 || 999.1026 ||
 * +10 || 999.7026 ||
 * **+4** || **999.9720** ||
 * 0 || 999.8395 ||
 * −10 || 998.117 ||
 * −20 || 993.547 ||
 * −30 || 983.854 ||
 * The values below 0 °C refer to [|supercooled] water. ||

Density of saltwater and ice
   [|WOA] surface density. 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 (see [|here] for explanation) and lowers the temperature of the density maximum of water to the freezing point. This 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. In cold countries, when the temperature of fresh water reaches 4 °C, the layers of water near the top in contact with cold air continue to lose heat energy and their temperature falls below 4 °C. On cooling below 4 °C, these layers do not sink but may rise up as fresh water has a maximum density at 4 °C. (Refer: Polarity and hydrogen bonding) Due to this, the layer of water at 4 °C remains at the bottom and above this layers of water 3 °C, 2 °C, 1 °C and 0 °C are formed. Since ice is a poor conductor of heat, it does not absorb heat energy from the water beneath the layer of ice which prevents the water freezing. Thus, aquatic creatures survive in such places.[// [|citation needed] //] 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 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 Poles, leading to a global system of currents called the [|thermohaline circulation]. One potential consequence of [|global warming] is that the loss of Arctic and Antarctic ice could result in the loss of these currents as well, which could have unforeseeable consequences on near and distant climates.

Miscibility and condensation
Red line shows saturation Main article: [|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 [|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] // [|[25]] 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 the humidity is increased at room temperature, for example, 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 then condenses out as minute water droplets, commonly referred to 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, for example, by rising suddenly in an updraft

Vapor pressure
Vapor pressure diagrams of water Main article: [|Vapor pressure of water]
 * ~ Temperature ||||||||||~ Pressure [|[27]] ||
 * ~ °C ||~ K ||~ °F ||~ Pa ||~ atm ||~ torr ||~ in Hg ||~ psi ||
 * 0 || 273 || 32 || 611 || 0.00603 || 4.58 || 0.180 || 0.0886 ||
 * 5 || 278 || 41 || 872 || 0.00861 || 6.54 || 0.257 || 0.1265 ||
 * 10 || 283 || 50 || 1,228 || 0.01212 || 9.21 || 0.363 || 0.1781 ||
 * 12 || 285 || 54 || 1,403 || 0.01385 || 10.52 || 0.414 || 0.2034 ||
 * 14 || 287 || 57 || 1,599 || 0.01578 || 11.99 || 0.472 || 0.2318 ||
 * 16 || 289 || 61 || 1,817 || 0.01793 || 13.63 || 0.537 || 0.2636 ||
 * 17 || 290 || 63 || 1,937 || 0.01912 || 14.53 || 0.572 || 0.2810 ||
 * 18 || 291 || 64 || 2,064 || 0.02037 || 15.48 || 0.609 || 0.2993 ||
 * 19 || 292 || 66 || 2,197 || 0.02168 || 16.48 || 0.649 || 0.3187 ||
 * 20 || 293 || 68 || 2,338 || 0.02307 || 17.54 || 0.691 || 0.3392 ||
 * 21 || 294 || 70 || 2,486 || 0.02453 || 18.65 || 0.734 || 0.3606 ||
 * 22 || 295 || 72 || 2,644 || 0.02609 || 19.83 || 0.781 || 0.3834 ||
 * 23 || 296 || 73 || 2,809 || 0.02772 || 21.07 || 0.830 || 0.4074 ||
 * 24 || 297 || 75 || 2,984 || 0.02945 || 22.38 || 0.881 || 0.4328 ||
 * 25 || 298 || 77 || 3,168 || 0.03127 || 23.76 || 0.935 || 0.4594 ||

Compressibility
The [|compressibility] of water is a function of pressure and temperature. At 0 °C, at the limit of zero pressure, the compressibility is 5.1×10−10 Pa−1. [|[28]] At the zero-pressure limit, the compressibility reaches a minimum of 4.4×10−10 Pa−1 around 45 °C before increasing again with increasing temperature. As the pressure is increased, the compressibility decreases, being3.9×10−10 Pa−1 at 0 °C and 100 MPa. The [|bulk modulus] of water is 2.2 GPa. [|[29]] The low compressibility of non-gases, and of water in particular, leads to their often being assumed as incompressible. The low compressibility of water means that even in the deep [|oceans] at 4 km depth, where pressures are 40 MPa, there is only a 1.8% decrease in volume

Triple point
The various triple points of water [|[30]] ||~ Phases in stable equilibrium 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.    Water [|phase diagram] : //Y//-axis = Pressure in pascals (10n); //X//-axis = temperature in kelvins; S = solid; L = liquid; V = vapor; CP = critical point; TP = [|triple point of water] 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 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 vapor 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
 * ~ Pressure ||~ Temperature ||
 * liquid water, [|ice Ih], and water vapor || 611.73 Pa || 273.16 K (0.01 °C) ||
 * liquid water, ice Ih, and [|ice III] || 209.9 MPa || 251 K (−22 °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 ||

Electrical conductivity
Pure water containing no exogenous ions is an excellent [|insulator], but not even "deionized" water is completely free of ions. Water undergoes [|auto-ionization] in the liquid state, when two water molecules form one [|hydroxide] anion (OH−) and one [|hydronium] cation ( H 3 O + ). Because water is such a good solvent, it almost always has some [|solute] dissolved in it, often a [|salt]. If water has even a tiny amount of such an impurity, then it can conduct electricity far more readily. 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 ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in otherwise ultra-pure water begins to noticeably lower its resistivity by up to several [|kΩ] ·m.[// [|citation needed] //] 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. In ice, the primary charge carriers are [|protons] (see [|proton conductor] ). [|[33]]

Electrolysis
Main article: [|Electrolysis of water] 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 attracted 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.

Polarity and hydrogen bonding
   Model of [|hydrogen bonds] (1) 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] meaning two poles. The oxygen end is partially negative and the hydrogen end is partially positive, because of this the direction of the [|dipole moment] points towards the oxygen. 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. [|[34]]

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 behavior 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, resulting in the anomalous decrease of density when cooled below 4 °C.

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.

Surface tension
Main article: [|Surface tension]    This [|paper clip] is under the water level, which has risen gently and smoothly. Surface tension prevents the clip from submerging and the water from overflowing the glass edges. Temperature dependence of the surface tension of pure water Water has a high [|surface tension] of 72.8 mN/m at [|room temperature], caused by the strong cohesion between water molecules, the highest of the non-metallic liquids. This can be seen when small quantities of water are placed onto a [|sorption] -free (non-adsorbent and non-absorbent) surface, such as [|polyethylene] 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.[// [|citation needed] //] 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.

[ [|edit] ] Capillary action
Main article: [|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 adhesive properties maintain the water attachment to the xylem and prevent tension rupture caused by [|transpiration pull].

Water as a solvent
Main article: [|aqueous solution]    Presence of [|colloidal] [|calcium carbonate] from high concentrations of dissolved [|lime] turns the water of [|Havasu Falls] 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 ( [|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 non-polar substances such as fats and oils are not. Non-polar 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 non-polar 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 in acid-base reactions
Chemically, water is [|amphoteric] : it can act as either an [|acid] or a [|base] in chemical reactions. According to the [|Brønsted-Lowry] definition, an acid is defined as a species which donates a proton (a 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 instance, water receives an H+ ion from HCl when [|hydrochloric acid] is formed: HCl (acid) + H2O (base) H3O+ + Cl− In the reaction with [|ammonia], NH3 , water donates a H+ ion, and is thus acting as an acid: NH3 (base) + H2O (acid) NH   + 4 + OH− Because the oxygen atom in water has two [|lone pairs], water often acts as a [|Lewis base] , or electron pair donor, in reactions with [|Lewis acids] , although it can also react with Lewis bases, forming hydrogen bonds between the electron pair donors and the hydrogen atoms of water. [|HSAB theory] describes water as both a weak hard acid and a weak hard base, meaning that it reacts preferentially with other hard species: H+ (Lewis acid) + H2O (Lewis base) → H3O+Fe3+ (Lewis acid) + H2O (Lewis base) → Fe(H2O) 3+ 6   Cl− (Lewis base) + H2O (Lewis acid) → Cl(H2O) − 6  When a salt of a weak acid or of a weak base is dissolved in water, water can partially [|hydrolyze] the salt, producing the corresponding base or acid, which gives aqueous solutions of [|soap] and [|baking soda] their basic pH: Na2CO3 + H2O NaOH + NaHCO3