View source for Properties of water

{{short description|Physical and chemical properties of pure water}}
{{redirect|HOH}}
{{broader|Water}}
<section begin=Chembox /><!-- TEMPLATES BEGIN -->
{{chembox
| verifiedrevid  = 477162647
| Name           = Water
| ImageFile      = 
| ImageFile1     = H2O 2D labelled.svg
| ImageSize1     = 150px
| ImageName1     = The water molecule has this basic geometric structure
| ImageClass1    = skin-invert-image
| ImageFileL1    = Water-3D-balls-A.png
| ImageSizeL1    = 100px
| ImageNameL1    = Ball-and-stick model of a water molecule
| ImageFileR1    = Water molecule 3D.svg
| ImageSizeR1    = 100px
| ImageNameR1    = Space filling model of a water molecule
| ImageCaptionR1 = {{legend|red|Oxygen, O}}{{legend|white|Hydrogen, H}}
| ImageFile2     = Water drop 001.jpg
| ImageSize2     = 264px
| ImageName2     = A drop of water falling towards water in a glass
| OtherNames     = {{ubl
|Hydrogen oxide
|Hydrogen hydroxide (HH or HOH)
|Hydroxylic acid
|[[Dihydrogen monoxide]] (DHMO) (parody name<ref>{{cite web |title=naming molecular compounds |url=http://www.iun.edu/~cpanhd/C101webnotes/chemical-nomenclature/bimolcmpds.html |website=www.iun.edu |access-date=1 October 2018 |quote=Sometimes these compounds have generic or common names (e.g., H2O is "water") and they also have systematic names (e.g., H2O, dihydrogen monoxide). |archive-date=24 September 2018 |archive-url=https://web.archive.org/web/20180924023206/http://www.iun.edu/~cpanhd/C101webnotes/chemical-nomenclature/bimolcmpds.html |url-status=dead }}</ref>)
|Dihydrogen oxide
|Hydric acid
|Hydrohydroxic acid
|Hydroxic acid
|Hydroxoic acid
|Hydrol<ref>{{cite encyclopedia| url = http://www.merriam-webster.com/dictionary/hydrol| title = Definition of Hydrol| dictionary = Merriam-Webster| access-date = 21 April 2019| archive-date = 13 August 2017| archive-url = https://web.archive.org/web/20170813221249/https://www.merriam-webster.com/dictionary/hydrol| url-status = live}}</ref>
|μ-Oxidodihydrogen
|κ<sup>1</sup>-Hydroxylhydrogen(0)
|Aqua
|Neutral liquid
|Oxygen dihydride (may be considered incorrect)
}}
| data page pagename = Water (data page)
| PIN            = Water
| SystematicName = Oxidane (not in common use){{sfn|Leigh|Favre|Metanomski|1998|p=99}}
| Section1       = {{Chembox Identifiers
| CASNo = 7732-18-5
| CASNo_Ref = {{cascite|correct|CAS}}
| Beilstein = 3587155
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 15377
| ChEMBL_Ref = {{ebicite|correct|EBI}}
| ChEMBL = 1098659
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 937
| DrugBank = DB09145
| EC_number = 231-791-2
| Gmelin = 117
| KEGG = C00001
| PubChem = 962 
| RTECS = ZC0110000
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = 059QF0KO0R
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/H2O/h1H2
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = XLYOFNOQVPJJNP-UHFFFAOYSA-N
| SMILES = O
}}
| Section2       = {{Chembox Properties
| Formula = {{chem|H|2|O}}
| MolarMass = 18.01528(33) g/mol
| Odor = Odorless
| Appearance = Almost colorless or white crystalline solid, almost colorless liquid, with a [[Color of water|hint of blue]], colorless gas<ref name="Braun_1993_612" />
| Density = {{ubl|Liquid (1 atm, [[Vienna Standard Mean Ocean Water|VSMOW]]):
|{{val|0.99984283|(84)|u=g/mL}} at {{val|0|u=degC}}<ref name=Tanaka>{{cite journal |last1=Tanaka |first1=M |last2=Girard |first2=G |last3=Davis |first3=R |last4=Peuto |first4=A |last5=Bignell |first5=N |title=Recommended table for the density of water between 0 C and 40 C based on recent experimental reports |journal=Metrologia |date=August 2001 |volume=38 |issue=4 |pages=301–309 |doi=10.1088/0026-1394/38/4/3}}</ref>
|{{val|0.99997495|(84)|u=g/mL}} at {{val|3.983035|(670)|u=degC}} (temperature of maximum density, often 4 °C)<ref name=Tanaka/>
|{{val|0.99704702|(83)|u=g/mL}} at {{val|25|u=degC}}<ref name=Tanaka/>
|{{val|0.96188791|(96)|u=g/mL}} at {{val|95|u=degC}}<ref>{{cite book|first1=Eric W.|last1=Lemmon|first2=Ian H.|last2=Bell|first3=Marcia L.|last3=Huber|first4=Mark O.|last4=McLinden|chapter=Thermophysical Properties of Fluid Systems|title=NIST Chemistry WebBook, NIST Standard Reference Database Number 69|editor-first1=P.J.|editor-last1=Linstrom|editor-first2=W.G.|editor-last2=Mallard|publisher=National Institute of Standards and Technology|doi=10.18434/T4D303|access-date=17 October 2023|url=https://webbook.nist.gov/cgi/fluid.cgi?Action=Load&ID=C7732185&Type=IsoBar&Digits=12&P=1&THigh=100&TLow=90&TInc=1&RefState=DEF&TUnit=C&PUnit=atm&DUnit=g%2Fml&HUnit=kJ%2Fmol&WUnit=m%2Fs&VisUnit=uPa*s&STUnit=N%2Fm|archive-date=23 October 2023|archive-url=https://web.archive.org/web/20231023083019/https://webbook.nist.gov/cgi/fluid.cgi?Action=Load&ID=C7732185&Type=IsoBar&Digits=12&P=1&THigh=100&TLow=90&TInc=1&RefState=DEF&TUnit=C&PUnit=atm&DUnit=g%2Fml&HUnit=kJ%2Fmol&WUnit=m%2Fs&VisUnit=uPa*s&STUnit=N%2Fm|url-status=live}}</ref>
|Solid:
|0.9167 g/mL at 0 °C{{sfn|Lide|2003|loc = Properties of Ice and Supercooled Water in Section 6}}
}}
| MeltingPtC = 0.00
| MeltingPt_notes ={{Efn|[[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).  Other isotopic compositions melt or boil at slightly different temperatures.|name = VSMOW}}
| BoilingPtC = 99.98
| BoilingPt_ref =<ref name="nist" />{{Efn|name=VSMOW}}
| pKa = 13.995{{sfn|Lide|2003|loc=Chapter 8: Dissociation Constants of Inorganic Acids and Bases}}{{sfn|Weingärtner et al.|2016|p = 13}}{{efn|
A commonly quoted value of 15.7 used mainly in organic chemistry for the pK<sub>a</sub> of water is incorrect.<ref>{{cite web| title = What is the pKa of Water| url = http://chemwiki.ucdavis.edu/Core/Organic_Chemistry/Fundamentals/What_is_the_pKa_of_water%3F| publisher = [[University of California, Davis]]| date = 2015-08-09| access-date = 2016-04-09| archive-date = 2016-02-14| archive-url = https://web.archive.org/web/20160214222524/http://chemwiki.ucdavis.edu/Core/Organic_Chemistry/Fundamentals/What_is_the_pKa_of_water%3F| url-status = live}}</ref><ref>{{cite journal|last1=Silverstein|first1=Todd P.|last2=Heller|first2=Stephen T.|title=pKa Values in the Undergraduate Curriculum: What Is the Real pKa of Water?|journal=Journal of Chemical Education|date=17 April 2017|volume=94|issue=6|pages=690–695|doi=10.1021/acs.jchemed.6b00623|bibcode=2017JChEd..94..690S}}</ref>}}
| pKb = 13.995
| ConjugateAcid = [[Hydronium]] H<sub>3</sub>O<sup>+</sup> (pK<sub>a</sub> = 0)
| ConjugateBase = [[Hydroxide]] OH<sup>–</sup> (pK<sub>b</sub> = 0)
| Viscosity = 0.890 mPa·s (0.890 [[Poise (unit)|cP]]){{sfn|Lide|2003|loc=6.186}}
| RefractIndex = 1.3330 (20&nbsp;°C){{sfn|Lide|2003|loc=8—Concentrative Properties of Aqueous Solutions: Density, Refractive Index, Freezing Point Depression, and Viscosity}}
| VaporPressure = {{convert|3.1690|kPa|atm|abbr = out| disp = or}} at 25&nbsp;°C{{sfn|Lide|2003|loc= Vapor Pressure of Water From 0 to 370&nbsp;°C in Sec. 6}}

| SolubleOther =  Poorly soluble in [[haloalkane]]s, [[aliphatic]] and [[aromatic]] hydrocarbons, [[ether]]s.<ref name=chemister>{{cite web| title = Properties of substance: water| url = http://chemister.ru/Database/properties-en.php?dbid=1&id=1| first = Kiper Ruslan| last = Anatolievich| access-date = 2014-06-01| archive-date = 2014-06-02| archive-url = https://web.archive.org/web/20140602200111/http://chemister.ru/Database/properties-en.php?dbid=1&id=1| url-status = live}}</ref> {{br}} Improved solubility in [[carboxylate]]s, [[alcohols]], [[ketones]], [[amines]]. {{br}} Miscible with [[methanol]], [[ethanol]], [[propanol]], [[isopropanol]], [[acetone]], [[glycerol]], [[1,4-dioxane]], [[tetrahydrofuran]], [[sulfolane]], [[acetaldehyde]], [[dimethylformamide]], [[dimethoxyethane]], [[dimethyl sulfoxide]], [[acetonitrile]]. {{br}} Partially miscible with [[diethyl ether]], [[methyl ethyl ketone]], [[dichloromethane]], [[ethyl acetate]], [[bromine]].
| ThermalConductivity = 0.6065 W/(m·K)<ref>{{Cite journal|last1=Ramires|first1=Maria L. V.|last2=Castro|first2=Carlos A. Nieto de|last3=Nagasaka|first3=Yuchi|last4=Nagashima|first4=Akira|last5=Assael|first5=Marc J.|last6=Wakeham|first6=William A.|date=1995-05-01|title=Standard Reference Data for the Thermal Conductivity of Water|journal=Journal of Physical and Chemical Reference Data|volume=24|issue=3|pages=1377–1381|doi=10.1063/1.555963|issn=0047-2689|bibcode=1995JPCRD..24.1377R}}</ref>
}}
| Section3       = {{Chembox Structure
| MolShape = [[bent (chemistry)|Bent]]
| CrystalStruct = [[ice|Hexagonal]]
| Dipole = 1.8546 [[Debye|D]]{{sfn|Lide|2003|loc=9—Dipole Moments}}
| PointGroup = C<sub>2v</sub>
}}
| Section4       = {{Chembox Thermochemistry
| DeltaHf = −285.83 ± 0.04&nbsp;kJ/mol<ref name=chemister /><ref name=nist>{{nist|name=Water|id=C7732185|access-date=2014-06-01|mask=FFFF|units=SI}}</ref>
| DeltaGf = −237.24&nbsp;kJ/mol<ref name=chemister />
| Entropy = 69.95 ± 0.03&nbsp;J/(mol·K)<ref name=nist />
| HeatCapacity = 75.385 ± 0.05&nbsp;J/(mol·K)<ref name=nist />
}}
| Section5       = 
| Section6       = 
| Section7       = {{Chembox Hazards
| MainHazards = [[Drowning]]<br />[[Avalanche]] (as snow)<br />[[Water intoxication]]
| NFPA-H = 0
| NFPA-F = 0
| NFPA-R = 0
| FlashPt= Non-flammable
| GHSPictograms = 
| GHSSignalWord =
| HPhrases = 
| PPhrases = 
| GHS_ref =<ref>GHS: [https://pubchem.ncbi.nlm.nih.gov/compound/962 PubChem 962] {{Webarchive|url=https://web.archive.org/web/20230728173927/https://pubchem.ncbi.nlm.nih.gov/compound/962 |date=2023-07-28 }}</ref>
| ExternalSDS = [http://www.labchem.com/tools/msds/msds/LC26750.pdf SDS]
}}
| Section9       = {{Chembox Related
| OtherAnions = {{ubl|[[Hydrogen sulfide]]|[[Hydrogen selenide]]|[[Hydrogen telluride]]|[[Hydrogen polonide]]|[[Hydrogen peroxide]]}}
| OtherFunction_label = [[solvent]]s
| OtherFunction = {{ubl|[[Acetone]]|[[Ethanol]]|[[Methanol]]|[[Hydrogen fluoride]]|[[Ammonia]]}}
}}
}}
<section end=Chembox /><!-- TEMPLATES END -->
'''Water''' ('''{{chem2|H2O}}''') is a [[Chemical polarity|polar]] inorganic compound that is at [[room temperature]] a tasteless and odorless [[liquid]], which is nearly colorless apart from [[Color of water|an inherent hint of blue]].<!--please read the article before considering removing it.--> It is by far the most studied chemical compound{{sfn|Greenwood|Earnshaw|1997|page=620}} and is described as the "universal [[solvent]]"<ref>{{cite web |url=https://www.usgs.gov/special-topic/water-science-school/science/water-universal-solvent?qt-science_center_objects=0#qt-science_center_objects |title=Water, the Universal Solvent |date=October 22, 2019 |department=U.S. Department of the Interior |website=usgs.gov |publisher=USGS |location=United States of America |language=en |type=website |access-date=December 15, 2020 |archive-date=December 1, 2021 |archive-url=https://web.archive.org/web/20211201104439/https://www.usgs.gov/special-topic/water-science-school/science/water-universal-solvent?qt-science_center_objects=0#qt-science_center_objects |url-status=live }}</ref> and the "solvent of life".{{sfn|Reece et al.|2013|p = 48}} It is the most abundant substance on the surface of [[Earth]]{{sfn|Weingärtner et al.|2016|p = 2}} and the only common substance to exist as a [[ice|solid]], liquid, and [[water vapor|gas]] on Earth's surface.{{sfn|Reece et al.|2013|p=44}} It is also the third most abundant molecule in the universe (behind [[Hydrogen|molecular hydrogen]] and [[carbon monoxide]]).{{sfn|Weingärtner et al.|2016|p = 2}}

Water molecules form [[hydrogen bonds]] with each other and are strongly polar. This polarity allows it to dissociate [[ions]] in salts and bond to other polar substances such as alcohols and acids, thus dissolving them. Its hydrogen bonding causes its many unique properties, such as having a solid form less dense than its liquid form, a relatively high [[boiling point]] of 100&nbsp;°C for its [[molar mass]], and a high [[heat capacity]].

Water is [[amphoteric]], meaning that it can exhibit properties of an [[acid]] or a [[base (chemistry)|base]], depending on the pH of the solution that it is in; it readily produces both [[hydron (chemistry)|{{chem|H|+}}]] and [[hydroxide|{{chem|O|H|-}}]] ions.{{efn|name = H+ | {{chem|H|+}} represents {{chem|H|3|O|+|(H|2|O)|''n''}} and more complex ions that form.}} Related to its amphoteric character, it undergoes [[self-ionization of water|self-ionization]]. The product of the [[chemical activity|activities]], or approximately, the concentrations of {{chem|H|+}} and {{chem|O|H|-}} is a constant, so their respective concentrations are inversely proportional to each other.<ref name="IUPAC-2009">{{cite book|chapter-url = http://goldbook.iupac.org/html/A/A00532.html|publisher = IUPAC|language = en|doi = 10.1351/goldbook.A00532|title = IUPAC Compendium of Chemical Terminology|year = 2009|isbn = 978-0-9678550-9-7|chapter = Autoprotolysis constant|access-date = 2018-08-09|archive-date = 2019-04-29|archive-url = https://web.archive.org/web/20190429095612/http://goldbook.iupac.org/html/A/A00532.html|url-status = live}}</ref>

== Physical properties ==
{{See also|Water chemistry analysis}}
Water is the [[chemical substance]] with [[chemical formula]] {{chem|H|2|O}}; one [[molecule]] of water has two [[hydrogen]] [[atom]]s [[covalent]]ly [[chemical bond|bonded]] to a single [[oxygen]] atom.{{sfn|Campbell|Williamson|Heyden|2006}} Water is a tasteless, odorless liquid at [[standard conditions|ambient temperature and pressure]]. Liquid water has weak [[absorption band]]s at wavelengths of around 750&nbsp;nm which cause it to appear to have a blue color.<ref name = Braun_1993_612>{{Cite journal|last1=Braun|first1=Charles L.|last2=Smirnov|first2=Sergei N.|date=1993-08-01|title=Why is water blue?|journal=Journal of Chemical Education|volume=70|issue=8|pages=612|bibcode=1993JChEd..70..612B|doi=10.1021/ed070p612|issn=0021-9584|url=http://inside.mines.edu/fs_home/dwu/classes/CH353/study/Why%20is%20Water%20Blue.pdf|access-date=2018-08-09|archive-date=2019-12-01|archive-url=https://web.archive.org/web/20191201000418/http://inside.mines.edu/fs_home/dwu/classes/CH353/study/Why%20is%20Water%20Blue.pdf|url-status=live}}</ref> This can easily be observed in a water-filled bath or wash-basin whose lining is white. Large ice crystals, as in [[glacier]]s, also appear blue.

Under [[Standard temperature and pressure|standard conditions]], water is primarily a liquid, unlike other analogous hydrides of the [[Chalcogen|oxygen family]], which are generally gaseous. This unique property of water is due to [[hydrogen bonding]]. The molecules of water are constantly moving concerning each other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200 femtoseconds (2&nbsp;×&nbsp;10<sup>−13</sup> seconds).<ref>{{cite journal|title=Unified description of temperature-dependent hydrogen bond rearrangements in liquid water|last=Smith|first=Jared D.|author2=Christopher D. Cappa|author3=Kevin R. Wilson|author4=Ronald C. Cohen|author5=Phillip L. Geissler|author6=Richard J. Saykally|journal=Proc. Natl. Acad. Sci. USA|date=2005|volume=102|pmid=16179387|issue=40|pmc=1242322|pages=14171–14174|doi=10.1073/pnas.0506899102|bibcode=2005PNAS..10214171S|doi-access=free}}</ref> However, these bonds are strong enough to create many of the peculiar properties of water, some of which make it integral to life.

=== Water, ice, and vapor ===
Within the Earth's atmosphere and surface, the [[liquid phase]] is the most common and is the form that is generally denoted by the word "water". The [[solid|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 material|granular]] crystals, like [[snow]]. Aside from [[Ice Ih|common hexagonal crystalline ice]], other crystalline and amorphous [[Ice#Phases|phases of ice]] are known. The [[gaseous phase]] of water is known as [[water vapor]] (or [[steam]]). Visible steam and clouds are formed from minute droplets of water suspended in the air.

Water also forms a [[supercritical fluid]]. The [[critical temperature]] is 647 [[kelvin|K]] and the [[critical pressure]] is 22.064 [[pascal (unit)|MPa]]. In nature, this only rarely occurs in extremely hostile conditions. A likely example of naturally occurring supercritical water is in the hottest parts of deep water [[hydrothermal vents]], in which water is heated to the critical temperature by [[submarine volcano|volcanic]] [[plume (hydrodynamics)|plumes]] and the critical pressure is caused by the weight of the ocean at the extreme depths where the vents are located. This pressure is reached at a depth of about 2200 meters: much less than the mean depth of the ocean (3800 meters).<ref>{{Cite journal|last1=Deguchi|first1=Shigeru|last2=Tsujii|first2=Kaoru|date=2007-06-19|title=Supercritical water: a fascinating medium for soft matter|journal=Soft Matter|language=en|volume=3|issue=7|pages=797–803|doi=10.1039/b611584e|pmid=32900070|issn=1744-6848|bibcode=2007SMat....3..797D}}</ref>

==== Heat capacity and heats of vaporization and fusion ====
[[File:Heat of Vaporization Water.png|left|thumb|Heat of vaporization of water from melting to critical temperature]]
Water has a very high [[specific heat capacity]] of 4184&nbsp;J/(kg·K) at 20&nbsp;°C (4182&nbsp;J/(kg·K) at 25&nbsp;°C) —the second-highest among all the heteroatomic species (after [[ammonia]]), as well as a high [[heat of vaporization]] (40.65&nbsp;kJ/mol or 2257&nbsp;kJ/kg at the normal boiling point), 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. [[Ocean_heat_content|Most of the additional energy stored in the climate system since 1970 has accumulated in the oceans]].<ref>{{cite report|chapter=3: Observations: Ocean|last1=Rhein|first1=M.|author-link1=Monika Rhein|last2=Rintoul|first2=S.R.|year=2013|chapter-url=http://www.climatechange2013.org/images/report/WG1AR5_Chapter03_FINAL.pdf|page=257|quote=Ocean warming dominates the global energy change inventory.  Warming of the ocean accounts for about 93% of the increase in the Earth's energy inventory between 1971 and 2010 (high confidence), with the warming of the upper (0 to 700&nbsp;m) ocean accounting for about 64% of the total. Melting ice (including Arctic sea ice, ice sheets, and glaciers) and warming of the continents and atmosphere account for the remainder of the change in energy.|title=IPCC WGI AR5|access-date=2017-12-22|archive-date=2020-10-16|archive-url=https://web.archive.org/web/20201016161329/http://www.climatechange2013.org/images/report/WG1AR5_Chapter03_FINAL.pdf|url-status=live}}</ref>

The specific [[enthalpy of fusion]] (more commonly known as latent heat) of water is 333.55&nbsp;kJ/kg at 0&nbsp;°C: the same amount of energy is required to melt ice as to warm ice from −160&nbsp;°C up to its melting point or to heat the same amount of water by about 80&nbsp;°C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice of [[glacier]]s and [[drift ice]]. Before and since the advent of mechanical [[refrigeration]], ice was and still is in common use for retarding food spoilage.

The specific heat capacity of ice at −10&nbsp;°C is 2030&nbsp;J/(kg·K){{sfn|Lide|2003|loc=Chapter 6: Properties of Ice and Supercooled Water}} and the heat capacity of steam at 100&nbsp;°C is 2080&nbsp;J/(kg·K).{{sfn|Lide|2003|loc=6. Properties of Water and Steam as a Function of Temperature and Pressure}}

==== Density of water and ice ====
[[File:Density of ice and water (en).svg|thumb|Density of ice and water as a function of temperature|left]]The [[density]] of water is about {{convert|1|g/cm3|lb/ft3}}: this relationship was originally used to define the gram.<ref name="decree">{{cite web|url=http://smdsi.quartier-rural.org/histoire/18germ_3.htm|title=Decree on weights and measures|date=April 7, 1795|quote=''Gramme'', le poids absolu d'un volume d'eau pure égal au cube de la centième partie du mètre, et à la température de la glace fondante.|access-date=August 3, 2016|archive-date=February 25, 2013|archive-url=https://web.archive.org/web/20130225163152/http://smdsi.quartier-rural.org/histoire/18germ_3.htm|url-status=dead}}</ref> The density varies with temperature, but not linearly: as the temperature increases, the density rises to a peak at {{convert|3.98|°C|°F}} and then decreases;{{sfn|Greenwood|Earnshaw|1997|page=625}} the initial increase is unusual because most liquids undergo [[thermal expansion]] so that the density only decreases as a function of temperature. The increase observed for water from {{convert|0|°C|°F}} to {{convert|3.98|°C|°F}} and for a few other liquids{{efn|Negative thermal expansion is also observed in [[molten silica]].<ref>{{cite journal|url=http://www.engr.ucsb.edu/~shell/papers/2002_PRE_silica.pdf|last1=Shell|first1=Scott M.|last2=Debenedetti|first2=Pablo G.|last3=Panagiotopoulos|first3=Athanassios Z.|title=Molecular structural order and anomalies in liquid silica|journal=Phys. Rev. E|date=2002|doi=10.1103/PhysRevE.66.011202|pmid=12241346|volume=66|issue=1|page=011202|arxiv=cond-mat/0203383|bibcode=2002PhRvE..66a1202S|s2cid=6109212|access-date=2009-07-07|archive-url=https://web.archive.org/web/20160604062440/http://www.engr.ucsb.edu/~shell/papers/2002_PRE_silica.pdf|archive-date=2016-06-04|url-status=dead}}</ref> Also, fairly pure silicon has a negative coefficient of thermal expansion for temperatures between about 18 and 120 [[kelvin]]s.<ref>{{cite book | editor1-first = William C. | editor1-last = O'Mara | editor2-first = Robert B. | editor2-last = Herring | editor3-first = Lee P. | editor3-last = Hunt | title = Handbook of semiconductor silicon technology | place = Park Ridge, New Jersey | publisher = Noyes Publications | year = 1990 | page = 431 | chapter-url = https://books.google.com/books?id=COcVgAtqeKkC&pg=PA431 | isbn = 0-8155-1237-6 | access-date = 2010-07-11 | author = Bullis, W. Murray | chapter = Chapter 6 | archive-date = 2024-02-04 | archive-url = https://web.archive.org/web/20240204075306/https://books.google.com/books?id=COcVgAtqeKkC&pg=PA431 | url-status = live }}</ref>
}} is described as [[negative thermal expansion]]. Regular, [[Ice Ih|hexagonal ice]] is also less dense than liquid water—upon freezing, the density of water decreases by about 9%.<ref name="Perlman" />{{Efn|Other substances that expand on freezing are [[silicon]] ([[melting point]] of {{convert|1687|K|C F}}), [[gallium]] (melting point of {{convert|303|K|C F}}, [[germanium]] (melting point of {{convert|1211|K|C F}}), and [[bismuth]] (melting point of {{convert|545|K|C F}})}}
These peculiar effects are due to the highly directional bonding of water molecules via the hydrogen bonds: ice and liquid water at low temperature have comparatively low-density, low-energy open lattice structures. The breaking of hydrogen bonds on melting with increasing temperature in the range 0–4&nbsp;°C allows for a denser molecular packing in which some of the lattice cavities are filled by water molecules.{{sfn|Greenwood|Earnshaw|1997|page=625}}<ref name=Housecroft>{{cite book |last1=Housecroft |first1=Catherine E. |last2=Sharpe |first2=Alan G. |title=Inorganic Chemistry |date=2005 |publisher=Pearson Prentice-Hall |isbn=0130-39913-2 |pages=162–163 |edition=2nd}}</ref> Above 4&nbsp;°C, however, thermal expansion becomes the dominant effect,<ref name=Housecroft/> and water near the boiling point (100&nbsp;°C) is about 4% less dense than water at {{convert|4|°C|°F}}.<ref name="Perlman" />{{efn|(1-0.95865/1.00000) × 100% {{=}} 4.135%}}

Under increasing pressure, ice undergoes a number of transitions to other [[polymorphism (materials science)|polymorphs]] with higher density than liquid water, such as [[ice II]], [[ice III]], [[high-density amorphous ice]] (HDA), and [[very-high-density amorphous ice]] (VHDA).<ref>{{Cite journal|author1-link=Thomas Loerting|last1=Loerting|first1=Thomas|last2=Salzmann|first2=Christoph|last3=Kohl|first3=Ingrid|last4=Mayer|first4=Erwin|last5=Hallbrucker|first5=Andreas|date=2001-01-01|title=A second distinct structural "state" of high-density amorphous ice at 77 K and 1 bar|journal=Physical Chemistry Chemical Physics|language=en|volume=3|issue=24|pages=5355–5357|doi=10.1039/b108676f|issn=1463-9084|bibcode=2001PCCP....3.5355L}}</ref>{{sfn|Greenwood|Earnshaw|1997|page=624}}
[[File:Anomalous expansion of water Summer Winter.svg|thumb|left|Temperature distribution in a lake in summer and winter]]
The unusual density curve and lower density of ice than of water is essential for much of the life on earth—if water were most dense at the freezing point, then in winter the cooling at the surface would lead to convective mixing. Once 0&nbsp;°C are reached, the water body would freeze from the bottom up, and all life in it would be killed.<ref name="Perlman" /> Furthermore, given that water is a good thermal insulator (due to its&nbsp;heat capacity), some frozen lakes might not completely thaw in summer.<ref name="Perlman" /> As it is, the inversion of the density curve leads to a stable layering for surface temperatures below 4&nbsp;°C, and with the layer of ice that floats on top insulating the water below,{{sfn|Zumdahl|Zumdahl|2013|p = 493}} even e.g., [[Lake Baikal]] in central [[Siberia]] freezes only to about 1 m thickness in winter. In general, for deep enough lakes, the temperature at the bottom stays constant at about 4&nbsp;°C (39&nbsp;°F) throughout the year (see diagram).<ref name="Perlman">{{Cite web|url=http://water.usgs.gov/edu/density.html|title=Water Density|last=Perlman|first=Howard|website=The USGS Water Science School|access-date=2016-06-03|archive-date=2016-06-25|archive-url=https://web.archive.org/web/20160625143337/http://water.usgs.gov/edu/density.html|url-status=live}}</ref>

==== Density of saltwater and ice ====
[[File:Annual mean sea surface sea water density (World Ocean Atlas 2009).png|thumb|upright=1.4|[[World Ocean Atlas|WOA]] surface density]]
The density of saltwater depends on the dissolved salt content as well as the temperature. 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 1.9&nbsp;°C<ref name="National Ocean Service">{{Cite web|url=http://oceanservice.noaa.gov/facts/oceanfreeze.html|title=Can the ocean freeze?|website=National Ocean Service|publisher=National Oceanic and Atmospheric Administration|language=EN-US|access-date=2016-06-09|archive-date=2020-07-06|archive-url=https://web.archive.org/web/20200706161806/https://oceanservice.noaa.gov/facts/oceanfreeze.html|url-status=live}}</ref> (due to [[Freezing-point depression#Freezing-point depression of a solvent and a solute|freezing-point depression of a solvent containing a solute]]) and lowers the temperature of the density maximum of water to the former freezing point at 0&nbsp;°C. 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. So creatures that live at the bottom of cold oceans like the [[Arctic Ocean]] generally live in water 4&nbsp;°C colder than at the bottom of frozen-over [[fresh water]] lakes and rivers.

As the [[surface science|surface]] of saltwater begins to freeze (at −1.9&nbsp;°C<ref name="National Ocean Service" /> for normal salinity [[seawater]], 3.5%) the ice that forms is essentially salt-free, with about the same density as 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 produces essentially freshwater ice at −1.9&nbsp;°C<ref name="National Ocean Service" /> 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]].

==== Miscibility and condensation ====
[[File:Relative Humidity.png|thumb|right|Red line shows saturation]]
{{Main|Humidity}}
{{see also | List of water-miscible solvents}}

Water is [[miscible]] with many liquids, including [[ethanol]] in all proportions. Water and most [[oil]]s are immiscible, usually forming layers according to increasing density from the top. This can be predicted by comparing the [[Chemical polarity|polarity]]. Water being a relatively polar compound will tend to be miscible with liquids of high polarity such as ethanol and acetone, whereas compounds with low polarity will tend to be immiscible and poorly [[soluble]] such as with [[hydrocarbons]].

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's [[partial pressure]] is 2% of atmospheric pressure and the air is cooled from 25&nbsp;°C, starting at about 22&nbsp;°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 saturated gas or one with 100% relative humidity is when the vapor pressure of water in the air is at 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 the 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. Vapor pressure above 100% relative humidity is called supersaturated and can occur if the air is rapidly cooled, for example, by rising suddenly in an updraft.{{Efn|''[[Adiabatic cooling]]'' resulting from the [[ideal gas law]].}}

==== Vapour pressure ====
{{Main|Vapour pressure of water}}
[[File:Vapor Pressure of Water.png|thumb|center|upright 2|Vapor pressure diagrams of water]]

==== Compressibility ====
The [[compressibility]] of water is a function of pressure and temperature. At 0&nbsp;°C, at the limit of zero pressure, the compressibility is {{val|5.1|e=-10|u=Pa<sup>−1</sup>}}. At the zero-pressure limit, the compressibility reaches a minimum of {{val|4.4|e=-10|u=Pa<sup>−1</sup>}} around 45&nbsp;°C before increasing again with increasing temperature. As the pressure is increased, the compressibility decreases, being {{val|3.9|e=-10|u=Pa<sup>−1</sup>}} at 0&nbsp;°C and {{convert|100|MPa|bar}}.<ref>{{cite journal |last1=Fine|first1 = R.A.|last2=Millero|first2 = F.J. |date=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 |bibcode=1973JChPh..59.5529F}}</ref>

The [[bulk modulus]] of water is about 2.2&nbsp;GPa.<ref name = nave>{{cite web|title = Bulk Elastic Properties|author = Nave, R.|work = HyperPhysics|publisher = [[Georgia State University]]|url = http://hyperphysics.phy-astr.gsu.edu/hbase/hph.html|access-date = 2007-10-26|archive-date = 2007-10-28|archive-url = https://web.archive.org/web/20071028155517/http://hyperphysics.phy-astr.gsu.edu/hbase/hph.html|url-status = live}}</ref> The low compressibility of non-gasses, and of water in particular, leads to their often being assumed as incompressible. The low compressibility of water means that even in the deep [[ocean]]s at 4&nbsp;km depth, where pressures are 40&nbsp;MPa, there is only a 1.8% decrease in volume.<ref name = nave/>

The bulk modulus of water ice ranges from 11.3&nbsp;GPa at 0&nbsp;K up to 8.6&nbsp;GPa at 273&nbsp;K.<ref>{{cite journal|last1=Neumeier|first1=J.J.|date=2018|title=Elastic Constants, Bulk Modulus, and Compressibility of H<sub>2</sub>O Ice I''h'' for the Temperature Range 50 K-273 K|volume=47|issue=3|page=033101|journal=Journal of Physical and Chemical Reference Data|doi=10.1063/1.5030640|bibcode=2018JPCRD..47c3101N|s2cid=105357042|url=https://scholarworks.montana.edu/xmlui/handle/1/15308|access-date=2021-08-03|archive-date=2021-11-28|archive-url=https://web.archive.org/web/20211128214007/https://scholarworks.montana.edu/xmlui/handle/1/15308|url-status=live}}</ref> The large change in the compressibility of ice as a function of temperature is the result of its relatively large thermal expansion coefficient compared to other common solids.

==== Triple point ====
{{main|Triple point#Triple point of water}}
[[File:Phase_diagram_of_water.svg|thumb|right|375px|upright 2|The solid/liquid/vapor triple point of liquid water, [[ice Ih|ice I<sub>h</sub>]] and water vapor in the lower left portion of a water phase diagram.]] The [[temperature]] and [[pressure]] at which ordinary solid, liquid, and gaseous water coexist in equilibrium is a [[triple point]] of water. Since 1954, this point had been used to define the base unit of temperature, the [[kelvin]],<ref>{{cite web |title=Base unit definitions: Kelvin |url=https://physics.nist.gov/cuu/Units/kelvin.html |publisher=[[National Institute of Standards and Technology]] |access-date=9 August 2018 |archive-date=20 August 2018 |archive-url=https://web.archive.org/web/20180820203424/https://physics.nist.gov/cuu/Units/kelvin.html |url-status=live }}</ref>{{sfn|Weingärtner et al.|2016|p = 5}} but, [[2019 revision of the SI|starting in 2019]], the kelvin is now defined using the [[Boltzmann constant]], rather than the triple point of water.<ref>
{{cite conference
 |title=Proceedings of the 106th meeting
 |conference=International Committee for Weights and Measures
 |date=16–20 October 2017
 |url=https://www.bipm.org/utils/en/pdf/CIPM/CIPM2017-EN.pdf?page=23
 |conference-url=https://www.bipm.org/en/committees/cipm/meeting/106.html
 |location=Sèvres
 |access-date=19 November 2018
 |archive-date=27 January 2018
 |archive-url=https://web.archive.org/web/20180127202612/https://www.bipm.org/utils/en/pdf/CIPM/CIPM2017-EN.pdf?page=23
 |url-status=live
}}</ref>

Due to the existence of many [[Polymorphism (materials science)|polymorphs]] (forms) of ice, water has other triple points, which have either three polymorphs of ice or two polymorphs of ice and liquid in equilibrium.{{sfn|Weingärtner et al.|2016|p = 5}} [[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 name=Schleuter>{{Cite thesis |title=Impact of High Pressure — Low Temperature Processes on Cellular Materials Related to Foods |last=Schlüter |first=Oliver |publisher=[[Technische Universität Berlin]] |url=http://edocs.tu-berlin.de./diss/2003/schlueter_oliver.pdf |date=2003-07-28 |url-status=dead |archive-url=https://web.archive.org/web/20080309120350/http://edocs.tu-berlin.de/diss/2003/schlueter_oliver.pdf |archive-date=2008-03-09 }}</ref><ref>{{Cite book |last=Tammann |first=Gustav H.J.A |date=1925 |title=The States Of Aggregation |publisher=[[Constable & Robinson|Constable And Company]] }}</ref>{{sfn|Lewis|Rice|1922|}}
{| class="wikitable" 
|+The various triple points of water

!Phases in stable equilibrium
!Pressure
!Temperature
|-
| liquid water, [[ice Ih|ice I<sub>h</sub>]], and water vapor
| 611.657 Pa<ref>{{cite journal | doi=10.1256/qj.04.94 | volume=131 | issue=608 | title=Review of the vapour pressures of ice and supercooled water for atmospheric applications | year=2005 | journal=Quarterly Journal of the Royal Meteorological Society | pages=1539–1565 | last1=Murphy | first1=D. M. | bibcode=2005QJRMS.131.1539M | s2cid=122365938 | url=https://zenodo.org/record/1236243 | doi-access=free | access-date=2020-08-31 | archive-date=2020-08-18 | archive-url=https://web.archive.org/web/20200818105335/https://zenodo.org/record/1236243 | url-status=live }}</ref>
| 273.16 K (0.01&nbsp;°C)
|-
| liquid water, ice I<sub>h</sub>, and [[ice III]]
| 209.9 MPa
| 251 K (−22&nbsp;°C)
|-
| liquid water, ice III, and [[ice V]]
| 350.1 MPa
| −17.0&nbsp;°C
|-
| liquid water, ice V, and [[ice VI]]
| 632.4 MPa
| 0.16&nbsp;°C
|-
| ice I<sub>h</sub>, [[Ice II]], and ice III
| 213 MPa
| −35&nbsp;°C
|-
| ice II, ice III, and ice V
| 344 MPa
| −24&nbsp;°C
|-
| ice II, ice V, and ice VI
| 626 MPa
| −70&nbsp;°C
|}

==== Melting point ====
The melting point of ice is {{convert| 0|C|F 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 about {{convert|231 | K | C F}}.<ref>{{cite journal|date=2003|title=Supercooled and Glassy Water|url=http://polymer.bu.edu/hes/articles/ds03.pdf|journal=Physics Today|volume=56|issue=6|pages=40–46|bibcode=2003PhT....56f..40D|doi=10.1063/1.1595053|last1=Debenedetti|first1=P. G.|last2=Stanley|first2=H. E.|access-date=2011-11-22|archive-date=2018-11-01|archive-url=https://web.archive.org/web/20181101114735/http://polymer.bu.edu/hes/articles/ds03.pdf|url-status=live}}</ref> The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, by {{convert|0.0073|C-change|F-change}}/atm{{efn|
name = sharp | The source gives it as 0.0072{{formatnum:}}°C/atm. However the author defines an [[atmosphere (unit)|atmosphere]] as 1,000,000 dynes/cm<sup>2</sup> (a [[Bar (unit)|bar]]). Using the standard definition of atmosphere, 1,013,250 dynes/cm<sup>2</sup>, it works out to 0.0073{{formatnum:}}°C/atm.}} or about {{convert| 0.5|C-change| F-change}}/70 atm{{efn|name = sharp2|
Using the fact that 0.5/0.0073 {{=}} 68.5.}}{{sfn|Sharp|1988|p = 27}} as the stabilization energy of hydrogen bonding is exceeded by intermolecular repulsion, but as ice transforms into its polymorphs (see [[Ice#Phases|crystalline states of ice]]) above {{convert|209.9|MPa|atm|abbr=on}}, the melting point increases markedly [[Ice#At different pressures|with pressure]], i.e., reaching {{convert|355|K|C}} at {{convert|2.216|GPa|atm|abbr=on}} (triple point of [[Ice VII]]<ref name="IAPWS">{{cite web|url=http://www.iapws.org/relguide/MeltSub2011.pdf|title=Revised Release on the Pressure along the Melting and Sublimation Curves of Ordinary Water Substance|date=September 2011|publisher=[[IAPWS]]|access-date=2013-02-19|archive-date=2014-03-02|archive-url=https://web.archive.org/web/20140302101543/http://www.iapws.org/relguide/MeltSub2011.pdf|url-status=live}}</ref>).

=== Electrical properties ===
==== Electrical conductivity ====
Pure water containing no exogenous [[ion]]s is an excellent electronic [[Electrical insulation|insulator]], but not even "deionized" water is completely free of ions. Water undergoes [[self-ionization of water|autoionization]] in the liquid state when two water molecules form one hydroxide anion ({{chem|OH|−|}}) and one hydronium cation ({{chem|H|3|O|+}}). Because of autoionization, at ambient temperatures pure liquid water has a similar intrinsic charge carrier concentration to the semiconductor germanium and an intrinsic charge carrier concentration three orders of magnitude greater than the semiconductor silicon, hence, based on charge carrier concentration, water can not be considered to be a completely dielectric material or electrical insulator but to be a limited conductor of ionic charge.<ref>C. S. Fuller "Defect Interactions in Semiconductors" Chapter 5 pp. 192-221 in "Semiconductors" N. B. Hannay Ed. Reinhold, New York 1959</ref>

Because water is such a good solvent, it almost always has some [[solute]] dissolved in it, often a [[Salt (chemistry)|salt]]. If water has even a tiny amount of such an impurity, then the ions can carry charges back and forth, allowing the water to conduct electricity far more readily.

It is known that the theoretical maximum electrical resistivity for water is approximately 18.2 MΩ·cm (182 [[kilohm|kΩ]]·m) at 25&nbsp;°C.<ref name = CondResist/> This figure agrees well with what is typically seen on [[reverse osmosis]], [[ultrafiltration|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|date=October 2010}}

In pure water, sensitive equipment can detect a very slight [[electrical conductivity]] of 0.05501 ± 0.0001 [[Siemens (unit)|μS]]/[[Centimeter|cm]] at 25.00&nbsp;°C.<ref name = CondResist>{{Cite journal|last1=Light|first1=Truman S.|last2=Licht|first2=Stuart|last3=Bevilacqua|first3=Anthony C.|last4=Morash|first4=Kenneth R.|date=2005-01-01|title=The Fundamental Conductivity and Resistivity of Water|journal=Electrochemical and Solid-State Letters|language=en|volume=8|issue=1|pages=E16–E19|doi=10.1149/1.1836121|issn=1099-0062}}</ref> 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. In ice, the primary charge carriers are [[protons]] (see [[proton conductor]]).<ref>
{{cite web
 |author=Crofts, A.
 |date=1996
 |title=Lecture 12: Proton Conduction, Stoichiometry
 |url=http://www.life.uiuc.edu/crofts/bioph354/lect12.html
 |publisher=[[University of Illinois at Urbana-Champaign]]
 |access-date=2009-12-06
 |archive-date=2009-05-10
 |archive-url=https://web.archive.org/web/20090510045553/http://www.life.uiuc.edu/crofts/bioph354/lect12.html
 |url-status=dead
}}</ref> Ice was previously thought to have a small but measurable conductivity of 1{{E|-10}}&nbsp;S/cm, but this conductivity is now thought to be almost entirely from surface defects, and without those, ice is an insulator with an immeasurably small conductivity.{{sfn|Greenwood|Earnshaw|1997|page=625}}

=== Polarity and hydrogen bonding ===
{{See also|Chemical polarity}}
[[File:Water molecule - structure and dipole moment.png|thumb|Water molecule - structure and dipole moment]]
An important feature of water is its polar nature. The structure has a [[bent molecular geometry]] for the two hydrogens from the oxygen vertex. The oxygen atom also has two [[lone pairs]] of electrons. One effect usually ascribed to the lone pairs is that the H–O–H gas-phase bend angle is 104.48°,<ref>{{cite journal|date=1979|last1=Hoy|first1=AR|last2=Bunker|first2=PR|journal=Journal of Molecular Spectroscopy|volume=74|issue=1|pages=1–8|doi=10.1016/0022-2852(79)90019-5|title=A precise solution of the rotation bending Schrödinger equation for a triatomic molecule with application to the water molecule|bibcode = 1979JMoSp..74....1H }}</ref> which is smaller than the typical [[tetrahedral molecular geometry|tetrahedral]] angle of 109.47°. The lone pairs are closer to the oxygen atom than the electrons [[sigma bond]]ed to the hydrogens, so they require more space. The increased repulsion of the lone pairs forces the O–H bonds closer to each other. {{sfn|Zumdahl|Zumdahl|2013|p = 393}}

Another consequence of its [[structure]] is that water is a [[polar molecule]]. Due to the difference in [[electronegativity]], a [[bond dipole moment]] points from each H to the O, making the oxygen partially negative and each hydrogen partially positive. A large molecular [[Electric dipole moment|dipole]], points from a region between the two hydrogen atoms to the oxygen atom. The charge differences cause water molecules to aggregate (the relatively positive areas being attracted to the relatively negative areas). This attraction, [[hydrogen bond]]ing, explains many of the properties of water, such as its solvent properties.{{sfn|Campbell|Farrell|2007|pp = 37–38}}

Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for several of the water's physical properties. These properties include its relatively high [[melting point|melting]] and boiling point temperatures: more energy is required to break the hydrogen bonds between water molecules. In contrast, hydrogen sulfide ({{chem|H|2|S}}), has much weaker hydrogen bonding due to sulfur's lower electronegativity. {{chem|H|2|S}} is a gas at [[room temperature]], despite hydrogen sulfide having nearly twice the molar 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.

==== Cohesion and adhesion ====
{{anchor|Cohesion}}
[[File:Spider web Luc Viatour.jpg|thumb|right|[[Dew]] drops adhering to a [[spider web]]]]
Water molecules stay close to each other ([[cohesion (chemistry)|cohesion]]), due to the collective action of hydrogen bonds between water molecules. These hydrogen bonds are constantly breaking, with new bonds being formed with different water molecules; but at any given time in a sample of liquid water, a large portion of the molecules are held together by such bonds.{{sfn|Campbell|Reece|2009|p=47}}

Water also has high [[adhesion]] properties because of its polar nature. On 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.{{citation needed|date=March 2021}} 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—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.<ref>{{cite journal|last1=Chiavazzo|first1=Eliodoro|last2=Fasano|first2=Matteo|last3=Asinari|first3=Pietro|last4=Decuzzi|first4=Paolo|title=Scaling behaviour for the water transport in nanoconfined geometries|journal=Nature Communications|volume=5|pages=4565|doi=10.1038/ncomms4565|bibcode = 2014NatCo...5.4565C |year=2014|pmid=24699509|pmc=3988813}}</ref> They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.<ref>{{cite web|url=http://www.biophysics.org/education/parsegian.pdf |title=Physical Forces Organizing Biomolecules |work=Biophysical Society |url-status=unfit |archive-url=https://web.archive.org/web/20070807213655/http://www.biophysics.org/education/parsegian.pdf |archive-date=August 7, 2007 }}</ref>
{{clear}}
[[File:RainDrops1.jpg|thumb|Rain water flux from a canopy. Among the forces that govern drop formation: [[Surface tension]], [[Cohesion (chemistry)]], [[Van der Waals force]], [[Plateau–Rayleigh instability]].]]

==== Surface tension ====
[[File:Paper Clip Surface Tension 1 crop.jpg|thumb|left|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.]]
[[File:Temperature dependence surface tension of water.svg|thumb|right|Temperature dependence of the surface tension of pure water]]
Water has an unusually high [[surface tension]] of 71.99&nbsp;mN/m at 25&nbsp;°C{{sfn|Lide|2003|loc=Surface Tension of Common Liquids}} which is caused by the strength of the hydrogen bonding between water molecules.{{sfn|Reece et al.|2013|p = 46}} This allows insects to walk on water.{{sfn|Reece et al.|2013|p = 46}}

==== Capillary action ====
Because water has strong cohesive and adhesive forces, it exhibits capillary action.{{sfn|Zumdahl|Zumdahl|2013|pp = 458–459}} Strong cohesion from hydrogen bonding and adhesion allows trees to transport water more than 100 m upward.{{sfn|Reece et al.|2013|p = 46}}

==== Water as a solvent ====
{{Main|Aqueous solution}}
[[File:Havasu Falls 2 md.jpg|thumb|upright|Presence of [[colloid]]al [[calcium carbonate]] from high concentrations of dissolved [[Lime (mineral)|lime]] turns the water of [[Havasu Falls]] turquoise.]]
Water is an excellent [[solvent]] due to its high dielectric constant.{{sfn|Greenwood|Earnshaw|1997|p = 627}} Substances that mix well and dissolve in water are known as [[hydrophilic]] ("water-loving") substances, while those that do not mix well with water are known as [[hydrophobic]] ("water-fearing") substances.{{sfn|Zumdahl|Zumdahl|2013|p = 518}} The ability of a substance to dissolve in water is determined by whether or not the substance can match or better the strong [[dipole–dipole interaction|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)|precipitated out]] from the water. 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 ([[Solvation|hydration]]). The relatively small size of water molecules (~ 3 angstroms) 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 (chemistry)|alcohol]]s, and [[salt (chemistry)|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 non-polar molecules.

An example of an ionic solute is [[sodium chloride|table salt]]; the sodium chloride, NaCl, separates into {{chem|Na|+}} [[cation]]s and {{chem|Cl|-}} [[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.

==== Quantum tunneling ====
{{Main|Quantum tunneling of water}}
The [[quantum tunneling]] dynamics in water was reported as early as 1992. At that time it was known that there are motions which destroy and regenerate the weak [[hydrogen bond]] by internal rotations of the substituent water [[monomers]].<ref>{{Cite journal|last=Pugliano|first=N.|date=1992-11-01|title=Vibration-Rotation-Tunneling Dynamics in Small Water Clusters|language=en|publisher=Lawrence Berkeley Lab., CA (United States)|page=6|doi=10.2172/6642535|osti=6642535|url=https://digital.library.unt.edu/ark:/67531/metadc1190847/|website=UNT Digital Library|access-date=2019-07-05|archive-date=2020-08-01|archive-url=https://web.archive.org/web/20200801150019/https://digital.library.unt.edu/ark:/67531/metadc1190847/|url-status=live}}</ref> On 18 March 2016, it was reported that the hydrogen bond can be broken by quantum tunneling in the water [[hexamer]]. Unlike previously reported tunneling motions in water, this involved the concerted breaking of two hydrogen bonds.<ref>{{Cite journal|last1=Richardson|first1=Jeremy O.|last2=Pérez|first2=Cristóbal|last3=Lobsiger|first3=Simon|last4=Reid|first4=Adam A.|last5=Temelso|first5=Berhane|last6=Shields|first6=George C.|last7=Kisiel|first7=Zbigniew|last8=Wales|first8=David J.|last9=Pate|first9=Brooks H.|last10=Althorpe|first10=Stuart C.|date=2016-03-18|title=Concerted hydrogen-bond breaking by quantum tunneling in the water hexamer prism|journal=Science|language=en|volume=351|issue=6279|pages=1310–1313|doi=10.1126/science.aae0012|issn=0036-8075|pmid=26989250|bibcode = 2016Sci...351.1310R |doi-access=free}}</ref> Later in the same year, the discovery of the quantum tunneling of water molecules was reported.<ref>{{Cite journal|last=Kolesnikov|first=Alexander I.|date=2016-04-22|title=Quantum Tunneling of Water in Beryl: A New State of the Water Molecule|journal=Physical Review Letters|volume=116|issue=16|pages=167802|doi=10.1103/PhysRevLett.116.167802|pmid=27152824|bibcode=2016PhRvL.116p7802K|url=https://zenodo.org/record/1233837|doi-access=free|access-date=2019-09-08|archive-date=2020-11-18|archive-url=https://web.archive.org/web/20201118142111/https://zenodo.org/record/1233837|url-status=live}}</ref>

=== Electromagnetic absorption ===
{{Main|Electromagnetic absorption by water}}
Water is relatively transparent to [[visible light]], [[near ultraviolet]] light, and [[far-red]] light, but it absorbs most [[ultraviolet light]], [[infrared light]], and [[microwave]]s. Most [[Photoreceptor protein|photoreceptors]] and [[photosynthetic pigment]]s 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. Water's light blue color is caused by weak [[Absorption (electromagnetic radiation)|absorption]] in the red part of the [[visible spectrum]].<ref name = Braun_1993_612 /><ref>{{cite journal |last=Pope |author2=Fry |title=Absorption spectrum (380-700nm) of pure water. II. Integrating cavity measurements |journal=Applied Optics |volume=36 |issue=33 |pages=8710–23 |year=1996 |doi=10.1364/ao.36.008710 |pmid=18264420 |bibcode = 1997ApOpt..36.8710P |s2cid=11061625 }}</ref>

== Structure ==
[[File:3D model hydrogen bonds in water.svg|left|thumb|Model of [[hydrogen bond]]s (1) between molecules of water]]
A single water molecule can participate in a maximum of four [[hydrogen bond]]s because it can accept two bonds using the lone pairs on oxygen and donate two hydrogen atoms. Other molecules like [[hydrogen fluoride]], ammonia, and [[methanol]] can also form hydrogen bonds. However, they do not show anomalous [[thermodynamics|thermodynamic]], [[Kinematics|kinetic]], or structural properties like those observed in water because none of them can form four hydrogen bonds: either they cannot donate or accept hydrogen atoms, or there are [[steric]] effects in bulky residues. In water, intermolecular [[tetrahedral]] structures form due to the four hydrogen bonds, thereby forming an open structure and a three-dimensional bonding network, resulting in the anomalous decrease in density when cooled below 4&nbsp;°C. This repeated, constantly reorganizing unit defines a three-dimensional network extending throughout the liquid. This view is based upon neutron scattering studies and computer simulations, and it makes sense in the light of the unambiguously tetrahedral arrangement of water molecules in ice structures.

However, there is an alternative theory for the structure of water. In 2004, a controversial paper from [[Stockholm University]] suggested that water molecules in the liquid state typically bind not to four but only two others; thus forming chains and rings. The term "string theory of water" (which is not to be confused with the [[string theory]] of physics) was coined. These observations were based upon X-ray absorption spectroscopy that probed the local environment of individual oxygen atoms.<ref>{{cite journal|year=2008|title=Water—an enduring mystery|journal=Nature|volume=452|issue=7185|pages=291–292|doi=10.1038/452291a|pmid=18354466|author=Ball, Philip|bibcode = 2008Natur.452..291B |s2cid=4365814|doi-access=free}}</ref>

=== Molecular structure ===
{{See also|Molecular orbital diagram#Water}}
The repulsive effects of the two lone pairs on the oxygen atom cause water to have a [[Bent (chemistry)|bent]], not [[Linear molecular geometry|linear]], molecular structure,<ref>{{cite book |last1=Gonick |first1=Larry |last2=Criddle |first2=Craig |title=The cartoon guide to chemistry |url=https://archive.org/details/cartoonguidetoch00gonirich |url-access=registration |publisher=HarperResource |isbn=9780060936778 |page=[https://archive.org/details/cartoonguidetoch00gonirich/page/59 59] |edition=1st |language=en |chapter=Chapter 3 Togetherness |quote=Water, H<sub>2</sub>O, is similar. It has two electron pairs with nothing attached to them. They, too, must be taken into account. Molecules like NH<sub>3</sub> and H<sub>2</sub>O are called '''bent'''. |date=2005-05-03 }}</ref> allowing it to be polar. The hydrogen–oxygen–hydrogen angle is 104.45°, which is less than the 109.47° for ideal [[Orbital hybridisation#sp3|sp<sup>3</sup> hybridization]]. The [[valence bond theory]] explanation is that the oxygen atom's lone pairs are physically larger and therefore take up more space than the oxygen atom's bonds to the hydrogen atoms.<ref>{{cite book |author1=Theodore L. Brown |display-authors=etal |title=Chemistry : the central science |date=2015 |isbn=978-0-321-91041-7 |page=351 |edition=13 |chapter-url=https://etext.pearson.com/eplayer/pdfbook?bookid=44972&userid=25122654&smsuserid=96481183&languageid=1&roletypeid=2&scenario=3&sessionid=c7d854c5fe00acd0458e561b1f1664e9&invoketype=et1&bookserver=1&platform=1002&uid=20180820172418&ubd=20180820172418&ubsd=20180820172418&hsid=db3145001e7f51fea9e1f6b972ea59f7 |access-date=21 April 2019 |chapter=9.2 The Vsepr Model |publisher=Pearson |quote=Notice that the bond angles decrease as the number of nonbonding electron pairs increases. A bonding pair of electrons is attracted by both nuclei of the bonded atoms, but a nonbonding pair is attracted primarily by only one nucleus. Because a nonbonding pair experiences less nuclear attraction, its electron domain is spread out more in space than is the electron domain for a bonding pair (Figure 9.7). Nonbonding electron pairs, therefore, take up more space than bonding pairs; in essence, they act as large and fatter balloons in our analogy of Figure 9.5. As a result, ''electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and tend to compress bond angles'' |archive-date=4 February 2024 |archive-url=https://web.archive.org/web/20240204075312/https://login.pearson.com/v1/piapi/iesui/webchallenge?client_id=6oVj1comRlGtGamiOamwDArIbVgGcKrA&redirect_uri=https%3A%2F%2Fplus.pearson.com%2Fhome&nonce=123454321&prompt=login&login_success_url=https%3A%2F%2Fplus.pearson.com%2Fhome%3FiesCode%3DRI8UsOaQhB |url-status=live }}</ref> The [[molecular orbital theory]] explanation ([[Bent's rule]]) is that lowering the energy of the oxygen atom's nonbonding hybrid orbitals (by assigning them more s character and less p character) and correspondingly raising the energy of the oxygen atom's hybrid orbitals bonded to the hydrogen atoms (by assigning them more p character and less s character) has the net effect of lowering the energy of the occupied molecular orbitals because the energy of the oxygen atom's nonbonding hybrid orbitals contributes completely to the energy of the oxygen atom's lone pairs while the energy of the oxygen atom's other two hybrid orbitals contributes only partially to the energy of the bonding orbitals (the remainder of the contribution coming from the hydrogen atoms' 1s orbitals).

== Chemical properties ==
=== Self-ionization ===
{{Main|Self-ionization of water}}
In liquid water there is some [[Self-ionization of water|self-ionization]] giving [[hydronium]] ions and [[hydroxide]] ions.
: 2 {{chem|H|2|O}} {{eqm}} {{chem|H|3|O|+}} + {{chem|OH|-}}

The [[equilibrium constant]] for this reaction, known as the [[self ionization of water|ionic product]] of water, <math>K_{\rm w}=[{\rm{H_3O^+}}][{\rm{OH^-}}] </math>, has a value of about {{10^|-14}} at 25&nbsp;°C. At neutral [[pH]], the concentration of the [[hydroxide]] ion ({{chem|OH|-}}) equals that of the (solvated) hydrogen ion ({{chem|H|+}}), with a value close to 10<sup>−7</sup> mol L<sup>−1</sup> at 25&nbsp;°C.{{sfn|Boyd|2000|p = 105}} See [[Water (data page)#Self-ionization|data page]] for values at other temperatures.

The thermodynamic equilibrium constant is a quotient of [[thermodynamic activity|thermodynamic activities]] of all products and reactants including water:
: <math>K_{\rm eq} = \frac{a_{\rm{H_3O^+}} \cdot a_{\rm{OH^-}}}{a_{\rm{H_2O}}^2}</math>
However, for dilute solutions, the activity of a solute such as H<sub>3</sub>O<sup>+</sup> or OH<sup>−</sup> is approximated by its concentration, and the activity of the solvent H<sub>2</sub>O is approximated by 1, so that we obtain the simple ionic product <math>K_{\rm eq} \approx K_{\rm w}=[{\rm{H_3O^+}}][{\rm{OH^-}}]</math>

=== Geochemistry ===
The action of water on rock over long periods of time typically leads to [[weathering]] and [[water erosion]], physical processes that convert solid rocks and minerals into soil and sediment, but under some conditions chemical reactions with water occur as well, resulting in [[metasomatism]] or [[mineral hydration]], a type of chemical alteration of a rock which produces [[clay minerals]]. It also occurs when [[Portland cement]] hardens.

Water ice can form [[clathrate compounds]], known as [[clathrate hydrates]], with a variety of small molecules that can be embedded in its spacious crystal lattice. The most notable of these is [[methane clathrate]], 4 {{chem|CH|4|·23H|2|O}}, naturally found in large quantities on the ocean floor.

=== Acidity in nature ===
Rain is generally mildly acidic, with a pH between 5.2 and 5.8 if not having any acid stronger than carbon dioxide.{{sfn|Boyd|2000|p = 106}} If high amounts of [[nitrogen]] and [[sulfur]] oxides are present in the air, they too will dissolve into the cloud and raindrops, producing [[acid rain]].

== Isotopologues ==
Several [[isotope]]s of both hydrogen and oxygen exist, giving rise to several known [[isotopologue]]s of water. [[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 [[Hydrogen-1|protium]]. Only 155 [[Parts per million|ppm]] include [[deuterium]] ({{chem|2|H}} or D), a hydrogen isotope with one neutron, and fewer than 20 parts per [[quintillion]] include [[tritium]] ({{chem|3|H}} or T), which has two neutrons. Oxygen also has three stable isotopes, with {{chem|16|O}} present in 99.76%, {{chem|17|O}} in 0.04%, and {{chem|18|O}} in 0.2% of water molecules.<ref>{{cite web|url=http://www.iapws.org/relguide/fundam.pdf|title=Guideline on the Use of Fundamental Physical Constants and Basic Constants of Water|date=2001|work=[[IAPWS]]|access-date=2008-03-21|archive-date=2017-01-28|archive-url=https://web.archive.org/web/20170128215733/http://www.iapws.org/relguide/fundam.pdf|url-status=live}}</ref>

Deuterium oxide, {{chem|D|2|O}}, is also known as [[heavy water]] because of its higher density. It is used in [[nuclear reactor]]s as a [[neutron moderator]]. Tritium is [[radioactive]], decaying with a [[half-life]] of 4500 days; {{chem|T|H|O}} exists in nature only in minute quantities, being produced primarily via cosmic ray-induced nuclear reactions in the atmosphere. Water with one protium and one deuterium atom {{chem|H|D|O}} occur naturally in ordinary water in low concentrations (~0.03%) and {{chem|D|2|O}} in far lower amounts (0.000003%) and any such molecules are temporary as the atoms recombine.

The most notable physical differences between {{chem|H|2|O}} and {{chem|D|2|O}}, other than the simple difference in specific mass, involve properties that are affected by hydrogen bonding, such as freezing and boiling, and other kinetic effects. This is because the nucleus of deuterium is twice as heavy as protium, and this causes noticeable differences in bonding energies. The difference in boiling points allows the isotopologues to be separated. The [[self-diffusion]] coefficient of {{chem|H|2|O}} at 25&nbsp;°C is 23% higher than the value of {{chem|D|2|O}}.<ref name =Isotopeneffekt>{{cite journal|author=Hardy, Edme H. |author2=Zygar, Astrid |author3=Zeidler, Manfred D. |author4=Holz, Manfred |author5=Sacher, Frank D. |title=Isotope effect on the translational and rotational motion in liquid water and ammonia|journal=J. Chem. Phys.|volume=114|issue=7 |year= 2001|pages=3174–3181|doi=10.1063/1.1340584|bibcode = 2001JChPh.114.3174H }}</ref> 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 {{chem|D|2|O}}.

Consumption of pure isolated {{chem|D|2|O}} may affect biochemical processes—ingestion of large amounts impairs kidney and central nervous system function. Small quantities can be consumed without any ill-effects; humans are generally unaware of taste differences,<ref name="Sci35">
{{cite news
 | last1 = Urey
 | first1 = Harold C.
 | display-authors = 1
 | last2 = Failla
 | first2 = Gioacchino
 | date = 15 Mar 1935
 | title = Concerning the Taste of Heavy Water
 | periodical = Science
 | place = New York
 | publisher = The Science Press
 | volume = 81
 | issue = 2098
 | page = 273
 | doi=10.1126/science.81.2098.273-a| bibcode = 1935Sci....81..273U
}}</ref> but sometimes report a burning sensation<ref name="PopSci35">
{{cite news
 | date = Apr 1935
 | title = Experimenter Drinks 'Heavy Water' at $5,000 a Quart
 | periodical = Popular Science Monthly
 | place = New York
 | publisher = Popular Science Publishing
 | volume = 126
 | issue = 4
 | page = 17
 | url = https://books.google.com/books?id=MSoDAAAAMBAJ&pg=PA17
 | access-date = 7 Jan 2011
}}</ref> or sweet flavor.<ref name="PopSci37">
{{cite news
 | last = Müller
 | first = Grover C.
 | date = June 1937
 | title = Is 'Heavy Water' the Fountain of Youth?
 | periodical = Popular Science Monthly
 | place = New York
 | publisher = Popular Science Publishing
 | volume = 130
 | issue = 6
 | pages = 22–23
 | url = https://books.google.com/books?id=eiYDAAAAMBAJ&pg=PA22
 | access-date = 7 Jan 2011
}}</ref> Very large amounts of heavy water must be consumed for any toxicity to become apparent. Rats, however, are able to avoid heavy water by smell, and it is toxic to many animals.<ref name="PnB">{{Cite journal|last1=Miller| first1=Inglis J. Jr. |last2=Mooser|first2=Gregory|date=Jul 1979|title=Taste Responses to Deuterium Oxide|journal=Physiology & Behavior|volume=23|issue=1|pages=69–74|doi=10.1016/0031-9384(79)90124-0|pmid=515218|s2cid=39474797}}</ref>

[[Deuterium-depleted water|''Light water'']] refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard {{nowrap|155 ppm}} level.

== Occurrence ==
Water is the most abundant substance on Earth's surface and also the third most abundant molecule in the universe, after {{chem|H|2}} and {{chem|C|O}}.{{sfn|Weingärtner et al.|2016|p = 2}} 0.23 ppm of the earth's mass is water and 97.39% of the global water volume of 1.38{{e|9}} km<sup>3</sup> is found in the oceans.{{sfn|Weingärtner et al.|2016|p = 29}}

Water is far more prevalent in the outer Solar System, beyond a point called the [[Frost line (astrophysics)|frost line]], where the Sun's radiation is too weak to vaporize solid and liquid water (as well as other elements and chemical compounds with relatively low melting points, such as [[methane]] and [[ammonia]]). In the inner Solar System, planets, asteroids, and moons formed almost entirely of metals and silicates. Water has since been delivered to the inner Solar System via an as-yet unknown mechanism, theorized to be the impacts of asteroids or comets carrying water from the outer Solar System, where bodies contain much more water ice.<ref>{{Cite journal |last=Prockter |first=Louise M. |date=2005 |title=Ice in the Solar System |url=https://secwww.jhuapl.edu/techdigest/content/techdigest/pdf/V26-N02/26-02-Prockter.pdf |journal=Johns Hopkins APL Technical Digest |volume=26 |issue=2 |pages=175–188 |via=Applied Physics Laboratory |access-date=2023-04-11 |archive-date=2023-04-11 |archive-url=https://web.archive.org/web/20230411142817/https://secwww.jhuapl.edu/techdigest/content/techdigest/pdf/V26-N02/26-02-Prockter.pdf |url-status=live }}</ref> The difference between planetary bodies located inside and outside the frost line can be stark. Earth's mass is 0.000023% water, while [[Tethys (moon)|Tethys]], a moon of Saturn, is almost entirely made of water.<ref>{{Cite web |date=2006-02-28 |title=Planetologie und Fernerkundung |url=https://www.geo.fu-berlin.de/geol/fachrichtungen/planet/index.html |access-date=2023-04-11 |website=www.geo.fu-berlin.de |language=de |archive-date=2023-04-11 |archive-url=https://web.archive.org/web/20230411142821/https://www.geo.fu-berlin.de/geol/fachrichtungen/planet/index.html |url-status=live }}</ref>

== Reactions ==

=== Acid–base reactions ===
Water is [[amphoteric]]: it has the ability to act as either an [[acid]] or a [[base (chemistry)|base]] in chemical reactions.{{sfn|Zumdahl|Zumdahl|2013|p = 659}} According to the [[Brønsted-Lowry]] definition, an acid is a proton ({{Chem|H|+|link=Hydron (chemistry)}}) donor and a base is a proton acceptor.{{sfn|Zumdahl|Zumdahl|2013|p = 654}} When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid.{{sfn|Zumdahl|Zumdahl|2013|p = 654}} For instance, water receives an {{chem|
|H|+}} ion from HCl when [[hydrochloric acid]] is formed:
: {{underset|<br />(acid)|HCl}} + {{underset|<br />(base)|{{chem|H|2|O}}}} {{eqm}} {{chem|H|3|O|+}} + {{chem|Cl|-}}

In the reaction with [[ammonia]], {{chem|NH|3}}, water donates a {{chem|H|+}} ion, and is thus acting as an acid:
: {{underset|<br />(base)|{{chem|NH|3}}}} + {{underset|<br />(acid)|{{chem|H|2|O}}}} {{eqm}} {{chem|NH|4|+}} + {{chem|OH|-}}

Because the oxygen atom in water has two [[lone pair]]s, water often acts as a [[Lewis base]], or electron-pair donor, in reactions with [[Lewis acid]]s, 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:
: {{underset|<br />(Lewis acid)|{{chem|H|+}}}} + {{underset|<br />(Lewis base)|{{chem|H|2|O}}}} → {{chem|H|3|O|+}}
: {{underset|<br />(Lewis acid)|{{chem|Fe|3+}}}} + {{underset|<br />(Lewis base)|{{chem|H|2|O}}}} → {{chem||Fe({{chem|H|2|O}})|6|3+}}
: {{underset|<br />(Lewis base)|{{chem|Cl|-}}}} + {{underset|<br />(Lewis acid)|{{chem|H|2|O}}}} → {{chem||Cl({{chem|H|2|O}})|6|-}}

When a salt of a weak acid or of a weak base is dissolved in water, water can partially [[hydrolysis|hydrolyze]] the salt, producing the corresponding base or acid, which gives aqueous solutions of [[soap]] and [[baking soda]] their basic pH:
: {{chem|Na|2|CO|3}} + {{chem|H|2|O}} {{eqm}} NaOH + {{chem|NaHCO|3}}

=== Ligand chemistry ===
Water's Lewis base character makes it a common [[ligand]] in [[transition metal]] complexes, examples of which include [[metal aquo complex]]es such as {{chem|Fe(H|2|O)|6|2+}} to [[perrhenic acid]], which contains two water molecules coordinated to a [[rhenium]] center.  In solid [[water of crystallization|hydrates]],  water can be either a ligand or simply lodged in the framework, or both.  Thus, {{chem|link=Ferrous sulfate|Fe|SO|4|·7H|2|O}} consists of [Fe<sub>2</sub>(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> centers and one "lattice water". Water is typically a [[monodentate]] ligand, i.e., it forms only one bond with the central atom.{{sfn|Zumdahl|Zumdahl|2013|p = 984}}
[[File:H-bondingFeSO47aq.tif|thumb|342 px|left|Some hydrogen-bonding contacts in FeSO<sub>4</sub><sup>.</sup>7H<sub>2</sub>O. This [[metal aquo complex]] crystallizes with one molecule of "lattice" water, which interacts with the sulfate and with the [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> centers.]]

=== Organic chemistry ===
As a hard base, water reacts readily with organic [[carbocation]]s; for example in a [[hydration reaction]], a hydroxyl group ({{chem|O|H|-}}) and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an alcohol. When the addition of water to an organic molecule cleaves the molecule in two, [[hydrolysis]] is said to occur. Notable examples of hydrolysis are the [[saponification]] of fats and the [[digestion]] of proteins and [[polysaccharides]]. Water can also be a [[leaving group]] in [[SN2 reaction|S<sub>N</sub>2 substitution]] and [[Elimination reaction|E2 elimination]] reactions; the latter is then known as a [[dehydration reaction]].

=== Water in redox reactions ===
Water contains hydrogen in the [[oxidation state]] +1 and oxygen in the oxidation state −2.{{sfn|Zumdahl|Zumdahl|2013|p = 171}} It oxidizes chemicals such as [[hydrides]], [[Alkali metal|alkali]] metals, and some [[Alkaline earth metal|alkaline earth]] metals.<ref>{{Cite web|url=http://chemwiki.ucdavis.edu/Core/Inorganic_Chemistry/Descriptive_Chemistry/Main_Group_Reactions/Compounds/Hydrides|title=Hydrides|website=Chemwiki|date=2 October 2013|publisher=[[UC Davis]]|language=en-US|access-date=2016-06-25|archive-date=2016-06-22|archive-url=https://web.archive.org/web/20160622230032/http://chemwiki.ucdavis.edu/Core/Inorganic_Chemistry/Descriptive_Chemistry/Main_Group_Reactions/Compounds/Hydrides|url-status=live}}</ref>{{sfn|Zumdahl|Zumdahl|2013|pp = 932, 936}} One example of an alkali metal reacting with water is:{{sfn|Zumdahl|Zumdahl|2013|p = 338}}
: 2 Na + 2 {{chem|H|2|O}} → {{chem|H|2}} + 2 {{chem|Na|+}} + 2 {{chem|O|H|-}}

Some other reactive metals, such as [[aluminium]] and [[beryllium]], are oxidized by water as well, but their oxides adhere to the metal and form a [[Passivation (chemistry)|passive]] protective layer.{{sfn|Zumdahl|Zumdahl|2013|p = 862}} Note that the [[rusting]] of [[iron]] is a reaction between iron and oxygen{{sfn|Zumdahl|Zumdahl|2013|p = 981}} that is dissolved in water, not between iron and water.

[[Heterogeneous water oxidation|Water can be oxidized]] to emit oxygen gas, but very few oxidants react with water even if their reduction potential is greater than the potential of {{chem|O|2|/H|2|O}}. Almost all such reactions require a [[catalyst]].{{sfn|Charlot|2007|p = 275}} An example of the oxidation of water is:
: 4 {{chem|Ag|F|2}} + 2 {{chem|H|2|O}} → 4 AgF + 4 HF + {{chem|O|2}}

=== Electrolysis ===
{{Main|Electrolysis of water}}
Water can be split into its constituent elements, hydrogen, and oxygen, by passing an electric current through it.{{sfn|Zumdahl|Zumdahl|2013|p = 866}} This process is called electrolysis. The cathode half reaction is:
: 2 {{chem|H|+}} + 2 {{Subatomic particle|Electron}} → {{chem|H|2}}

The anode half reaction is:
: 2 {{chem|H|2|O}} → {{chem|O|2}} + 4 {{chem|H|+}} + 4 {{Subatomic particle|Electron}}

The gases produced bubble to the surface, where they can be collected or ignited with a flame above the water if this was the intention. The required potential for the electrolysis of pure water is 1.23 V at 25&nbsp;°C.{{sfn|Zumdahl|Zumdahl|2013|p = 866}} The operating potential is actually 1.48 V or higher in practical electrolysis.

== History ==
[[Henry Cavendish]] showed that water was composed of oxygen and hydrogen in 1781.{{sfn|Greenwood|Earnshaw|1997|page=601}} The first decomposition of water into hydrogen and oxygen, by [[electrolysis]], was done in 1800 by English chemist [[William Nicholson (chemist)|William Nicholson]] and [[Anthony Carlisle]].{{sfn|Greenwood|Earnshaw|1997|page=601}}<ref>{{Cite web|url=http://www.rsc.org/chemistryworld/Issues/2003/August/electrolysis.asp|title=Enterprise and electrolysis...|date=August 2003|website=Royal Society of Chemistry|access-date=2016-06-24|archive-date=2016-03-03|archive-url=https://web.archive.org/web/20160303201557/http://www.rsc.org/chemistryworld/Issues/2003/August/electrolysis.asp|url-status=live}}</ref> In 1805, [[Joseph Louis Gay-Lussac]] and [[Alexander von Humboldt]] showed that water is composed of two parts hydrogen and one part oxygen.<ref>{{Cite web|url=http://www.1902encyclopedia.com/G/GAY/joseph-louis-gay-lussac.html|title=Joseph Louis Gay-Lussac, French chemist (1778–1850)|website=1902 Encyclopedia|at=Footnote 122-1|access-date=2016-05-26|archive-date=2023-05-29|archive-url=https://web.archive.org/web/20230529053730/https://www.1902encyclopedia.com/G/GAY/joseph-louis-gay-lussac.html|url-status=live}}</ref>

[[Gilbert Newton Lewis]] isolated the first sample of pure [[heavy water]] in 1933.<ref>{{Cite journal| pages = 341| year = 1933 | doi = 10.1063/1.1749300| last2 =  MacDonald| first1 = G. N.| volume = 1| last1 = Lewis| journal = The Journal of Chemical Physics | first2 = R. T.| title = Concentration of H2 Isotope| issue = 6|bibcode = 1933JChPh...1..341L }}</ref>

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

== Nomenclature ==
The accepted [[IUPAC nomenclature of inorganic chemistry|IUPAC]] name of water is ''oxidane'' or simply ''water'',{{sfn|Leigh|Favre|Metanomski|1998|p = 34}} or its equivalent in different languages, although there are other systematic names which can be used to describe the molecule. Oxidane is only intended to be used as the name of the mononuclear [[parent hydride]] used for naming derivatives of water by [[Substitutive nomenclature|substituent nomenclature]].{{sfn|IUPAC|2005|p = 85}} These derivatives 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]].{{sfn|Leigh|Favre|Metanomski|1998|p = 99}}<ref>{{Cite web|url=https://pubchem.ncbi.nlm.nih.gov/compound/8894#section=IUPAC-Name|title=Tetrahydropyran|website=Pubchem|publisher=[[National Institutes of Health]]|access-date=2016-07-31|archive-date=2016-08-16|archive-url=https://web.archive.org/web/20160816175013/http://pubchem.ncbi.nlm.nih.gov/compound/8894#section=IUPAC-Name|url-status=live}}</ref>

The simplest systematic name of water is ''hydrogen oxide''. This is analogous to related compounds such as [[hydrogen peroxide]], [[hydrogen sulfide]], and [[deuterium oxide]] (heavy water). Using chemical nomenclature for [[chemical nomenclature#Type-I ionic binary compounds|type I ionic binary compounds]], water would take the name ''hydrogen monoxide''<!-- Why is this in here? Water is polar covalent, not ionic. Also, for covalent compounds with hydrogen first (and for ionic compounds), the numeric prefix is not necessary for either of the elements. -->,{{sfn|Leigh|Favre|Metanomski|1998|pp = 27–28}} but this is not among the names published by the [[International Union of Pure and Applied Chemistry]] (IUPAC).{{sfn|Leigh|Favre|Metanomski|1998|p = 34}}<!-- IUPAC does not give a limit to the names for substances and is not giving an exhaustive list of acceptable names, nor any guidance on what names should be used. On that very page it gives TRIhydrogen oxide --> Another name is ''dihydrogen monoxide'', which is a rarely used name of water, and mostly used in the [[dihydrogen monoxide parody]].

Other systematic names for water include ''hydroxic acid'', ''hydroxylic acid'', and ''hydrogen hydroxide'', using acid and base names.{{efn|Both acid and base names exist for water because it is [[amphoterism|amphoteric]] (able to react both as an acid or an alkali).}} None of these exotic names are used widely. The polarized form of the water molecule, {{chem|H|+|OH|−}}, is also called [[hydron (chemistry)|hydron]] hydroxide by IUPAC nomenclature.<ref>{{cite web|url=https://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=22247451&loc=ec_rcs|title=Compound Summary for CID 22247451|work=[[PubChem|Pubchem Compound Database]]|publisher=National Center for Biotechnology Information|access-date=2017-09-08|archive-date=2014-08-27|archive-url=https://web.archive.org/web/20140827083048/http://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=22247451&loc=ec_rcs|url-status=live}}</ref>

''Water substance'' is a rare term used for H<sub>2</sub>O when one does not wish to specify the phase of matter (liquid [[water]], [[water vapor]], some form of [[ice]], or a component in a mixture) though the term "water" is also used with this general meaning.

''Oxygen dihydride'' is another way of referring to water, but modern usage often restricts the term "[[hydride]]" to ionic compounds (which water is not).

== See also ==
{{Portal|Water}}
{{Colbegin|colwidth=20em}}
* [[Chemical bonding of water]]
* [[Dihydrogen monoxide parody]]
* [[Double distilled water]]
* [[Electromagnetic absorption by water]]
* [[Fluid dynamics]]
* [[Hard water]]
* [[Heavy water]]
* [[Hydrogen polyoxide]]
* [[Ice]]
* [[Optical properties of water and ice]]
* [[Steam]]
* [[Superheated water]]
* {{section link|Viscosity#Water|}}
* [[Water cluster]]
* [[Water (data page)]]
* [[Water dimer]]
* [[Water model]]
* [[Water thread experiment]]
{{Colend}}

== Footnotes ==
{{notelist|30em}}

== References ==
=== Notes ===
{{reflist|30em}}

=== Bibliography ===
{{refbegin|30em}}
* {{cite book|first1=Claude E.|last1=Boyd|title=Water Quality|pages=105–122|chapter = pH, Carbon Dioxide, and Alkalinity|publisher=Springer|location = Boston, Massachusetts|year = 2000|isbn=978-1461544852|doi=10.1007/978-1-4615-4485-2_7}}
* {{cite book|title=Biochemistry|publisher=Cengage Learning|date=2007|edition=6th|isbn=978-0-495-39041-1|url=https://books.google.com/books?id=NYa45_BxgukC&pg=PA37|last1=Campbell|first1=Mary K.|last2=Farrell|first2=Shawn O.}}
* {{cite book|title=Biology|url=https://archive.org/details/essentialbiology00camp_0|url-access=registration|publisher=Pearson|date=2009|edition=8th|isbn=978-0-8053-6844-4|last1=Campbell|first1=Neil A.|last2=Reece|first2=Jane B.}}
* {{cite book|last1 = Campbell|first1 = Neil A.|first2 = Brad|last2 = Williamson|first3 = Robin J.|last3 = Heyden|title = Biology: Exploring Life|publisher = Pearson Prentice Hall|date = 2006|location = Boston|url = http://www.phschool.com/el_marketing.html|isbn = 978-0-13-250882-7|access-date = 2008-11-19|archive-date = 2014-11-02|archive-url = https://web.archive.org/web/20141102041816/http://www.phschool.com/el_marketing.html|url-status = live}}
* {{cite book|url=https://books.google.com/books?id=Ml-AJ9YbnTIC|title=Qualitative Inorganic Analysis|last=Charlot|first=G.|isbn=978-1-4067-4789-8|date=2007|publisher=Read Books}}
* {{cite book|last1=Greenwood |first1=Norman N. |author-link1=Norman Greenwood |last2=Earnshaw |first2=Alan |year=1997 |title=Chemistry of the Elements |edition=2nd |publisher=[[Butterworth-Heinemann]] |isbn=978-0-08-037941-8}}
* {{cite book|author=International Union of Pure and Applied Chemistry|url=http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|title=Nomenclature of Inorganic Chemistry: IUPAC Recommendations 2005|date=2005|publisher=Royal Society of Chemistry|isbn=978-0-85404-438-2|access-date=2016-07-31|ref={{harvid|IUPAC|2005}}|author-link=International Union of Pure and Applied Chemistry|archive-date=2019-12-12|archive-url=https://web.archive.org/web/20191212225258/http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|url-status=live}}
* {{cite book |url=http://old.iupac.org/publications/books/principles/principles_of_nomenclature.pdf |title=Principles of chemical nomenclature: a guide to IUPAC recommendations |last1=Leigh |first1=G. J. |last2=Favre |first2=H. A |last3=Metanomski |first3=W. V. |date=1998 |publisher=Blackwell Science |location=Oxford |oclc=37341352 |isbn=978-0-86542-685-6 |url-status=dead |archive-url=https://web.archive.org/web/20110726171925/http://old.iupac.org/publications/books/principles/principles_of_nomenclature.pdf |archive-date=2011-07-26 }}
* {{cite book|title=A System of Physical Chemistry|first1=William C.M.|first2=James|date=1922|publisher=Longmans, Green and Co.|last1=Lewis|last2=Rice}}
* {{cite book|url=https://books.google.com/books?id=kTnxSi2B2FcC|title=CRC Handbook of Chemistry and Physics|edition=84th|last=Lide|first=David R.|date=2003|series=[[CRC Handbook]]|publisher=CRC Press|isbn=978-0849304842|language=en|access-date=2016-05-29|archive-date=2024-02-04|archive-url=https://web.archive.org/web/20240204075307/https://books.google.com/books?id=kTnxSi2B2FcC|url-status=live}}
* {{cite book|title=Campbell Biology|last1=Reece|first1=Jane B.|last2=Urry|first2=Lisa A.|last3=Cain|first3=Michael L.|last4=Wasserman|first4=Steven A.|last5=Minorsky|first5=Peter V.|last6=Jackson|first6=Robert B.|date=2013|publisher=Pearson|isbn=978-0321775658|edition=10th|location=Boston, Mass.|language=en|ref = {{harvid|Reece et al.|2013}}}}
* {{cite book|title=Organic Solvents Physical Properties and Methods of Purification|url=https://archive.org/details/organicsolventsp0002ridd|url-access=registration|last=Riddick|first=John|date=1970|publisher=Wiley-Interscience|isbn=978-0471927266|series=Techniques of Chemistry|language=en}}
* {{cite book|url=https://archive.org/details/livingiceunderst0000shar|url-access=registration|page=[https://archive.org/details/livingiceunderst0000shar/page/27 27]|title=Living Ice: Understanding Glaciers and Glaciation|last=Sharp|first=Robert Phillip|date=1988|publisher=Cambridge University Press|isbn=978-0-521-33009-1|author-link=Robert P. Sharp}}
* {{cite book|first1=Hermann|last1=Weingärtner|first2=Ilka|last2=Teermann|first3=Ulrich|last3=Borchers|first4=Peter|last4=Balsaa|first5=Holger V.|last5=Lutze|first6=Torsten C.|last6=Schmidt|first7=Ernst Ulrich|last7=Franck|first8=Gabriele|last8=Wiegand|first9=Nicolaus|last9=Dahmen|first10=Georg|last10=Schwedt|first11=Fritz H.|last11=Frimmel|first12=Birgit C.|last12=Gordalla|chapter=Water, 1. Properties, Analysis, and Hydrological Cycle|title = Ullmann's Encyclopedia of Industrial Chemistry|publisher=Wiley-VCH Verlag GmbH & Co. KGaA|isbn=978-3527306732|doi=10.1002/14356007.a28_001.pub3|year = 2016|ref = {{harvid|Weingärtner et al.|2016}}|title-link=Ullmann's Encyclopedia of Industrial Chemistry}}
* {{cite book|title=Chemistry|last1=Zumdahl|first1=Steven S.|last2=Zumdahl|first2=Susan A.|date=2013|publisher=[[Cengage Learning]]|isbn=978-1-13-361109-7|edition=9th}}
{{refend}}

== Further reading ==
* {{citation | author-first= A. | author-last= Ben-Naim | author-link= Arieh Ben-Naim | title= Molecular Theory of Water and Aqueous Solutions | year= 2011 | publisher= [[World Scientific]]}}

== External links ==
{{Commons|Water molecule}}
{{Commons category|Water diagrams}}
{{Wikiversity|Engineering thermodynamics/Steam tables}}
* {{cite web |url=http://water.usgs.gov/edu/waterproperties.html |title=Water Properties and Measurements |author=<!--Staff writer(s); no by-line.--> |date=May 2, 2016 |publisher=[[United States Geological Survey]] |access-date=August 31, 2016 }}
* [https://web.archive.org/web/20080309145159/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)
* [https://mslngr.de/tools/h2o/h2o_gui.html Online calculator using the IAPWS Supplementary Release on Properties of Liquid Water at 0.1 MPa, September 2008]
* {{cite book|last=Chaplin|first=Martin|title=Encyclopedia of Water|chapter=Structure and Properties of Water in its Various States|year=2019|pages=1–19|publisher=Wiley Online Library 2019|doi=10.1002/9781119300762.wsts0002|isbn=9781119300755|s2cid=213738895|chapter-url=https://doi.org/10.1002/9781119300762.wsts0002}}
* Calculation of [http://ddbonline.ddbst.de/AntoineCalculation/AntoineCalculationCGI.exe?component=Water vapor pressure], [http://ddbonline.ddbst.de/DIPPR105DensityCalculation/DIPPR105CalculationCGI.exe?component=Water liquid density], [http://ddbonline.ddbst.de/VogelCalculation/VogelCalculationCGI.exe?component=Water dynamic liquid viscosity], and [http://ddbonline.ddbst.de/DIPPR106SFTCalculation/DIPPR106SFTCalculationCGI.exe?component=Water surface tension] of water
* [http://www.linkingweatherandclimate.com/ocean/waterdensitycalc.php Water Density Calculator]
* [https://web.archive.org/web/20110806030851/http://www.nasa.gov/audience/foreducators/topnav/materials/listbytype/Why_Does_Ice_Float.html Why does ice float in my drink?], [[NASA]]

{{Water|expand}}
{{Navboxes|list=
{{Hydrogen compounds}}
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