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For other uses, see Water (disambiguation). "H2O" redirects here. For other uses, see H2O
(disambiguation).

Water

Oxygen, O

Hydrogen, H

Names

Preferred IUPAC name

Water

Systematic IUPAC name

Oxidane (not in common use)[3]

Other names

 Hydrogen oxide

 Hydrogen hydroxide (H2O or HOH)

 Hydroxylic acid

 Dihydrogen monoxide (DHMO) (parody

 name[1])

 Dihydrogen oxide

 Hydric acid

 Hydrohydroxic acid

 Hydroxic acid

 Hydroxoic acid

 Hydrol[2]

 μ-Oxidodihydrogen

 κ1-Hydroxylhydrogen(0)

 Aqua

 Neutral liquid

 Oxygen dihydride (may be considered

incorrect)

Identifiers

CAS  7732-18-5
Number

3D  Interactive image
model
(JSmol)

Beilstein 3587155
Referen
ce

ChEBI  CHEBI:15377

ChEMBL  ChEMBL1098659

ChemSp  937
ider

DrugBan  DB09145
k
ECHA 100.028.902
InfoCard

EC  231-791-2
Number

Gmelin 117
Referen
ce

KEGG  C00001

PubChe  962
m CID

RTECS  ZC0110000
number

UNII  059QF0KO0R

CompTo  DTXSID6026296
x
Dashbo
ard (EPA
)

show

InChI

show

SMILES

Properties

Chemica H
l 2O

formula

Molar 18.01528(33) g/mol

mass

Appeara Almost colorless or white crystalline solid,

nce almost colorless liquid, with a hint of blue,
colorless gas[4]

Odor Odorless

Density  Liquid (1 atm, VSMOW):

 0.99984283(84) g/mL at 0 °C[5]

 0.99997495(84) g/mL at 3.983035(6

70) °C (temperature of maximum
density, often 4 °C)[5]

 0.99704702(83) g/mL at 25 °C[5]

 0.96188791(96) g/mL at 95 °C[6]

 Solid:

 0.9167 g/mL at 0 °C[7]

Melting 0.00 °C (32.00 °F; 273.15 K) [b]

point

Boiling 99.98 °C (211.96 °F; 373.13 K)[17][b]

point

Solubilit Poorly soluble

y in haloalkanes, aliphatic and aromatic hydr
ocarbons, ethers.[8]
Improved solubility
in carboxylates, alcohols, ketones, amines.
Miscible
with methanol, ethanol, propanol, isoprop
anol, acetone, glycerol, 1,4-dioxane, tetrah
ydrofuran, sulfolane, acetaldehyde, dimeth
ylformamide, dimethoxyethane, dimethyl
sulfoxide, acetonitrile.
Partially miscible with diethyl ether, methyl
ethyl ketone, dichloromethane, ethyl
acetate, bromine.

Vapor 3.1690 kilopascals or 0.031276 atm at

pressure 25 °C[9]
Acidity ( 13.995[10][11][a]
pKa)

Basicity 13.995
(pKb)

Conjuga Hydronium H3O+ (pKa = 0)

te acid

Conjuga Hydroxide OH– (pKb = 0)

te base

Thermal 0.6065 W/(m·K)[14]

conducti
vity

Refracti 1.3330 (20 °C)[15]

ve
index (n
D)

Viscosit 0.890 mPa·s (0.890 cP)[16]

y

Structure

Crystal Hexagonal
structur
e

Point C2v
group

Molecul Bent
ar shape

Dipole 1.8546 D[18]

moment

Thermochemistry

Heat 75.385 ± 0.05 J/(mol·K)[17]

capacity
(C)

Std 69.95 ± 0.03 J/(mol·K)[17]

molar
entropy
(S⦵298)

Std −285.83 ± 0.04 kJ/mol[8][17]

enthalp
y of
formatio
n (ΔfH⦵298
)

Gibbs −237.24 kJ/mol[8]

free
energy (
ΔfG⦵)

Hazards

Occupational safety and health (OHS/OSH):

Main Drowning
hazard Avalanche (as snow)
s Water intoxication

NFPA
704 (fire
diamon
d)
0

Flash Non-flammable
point

Related compounds

Other a  Hydrogen sulfide

nions
  Hydrogen selenide

 Hydrogen telluride

 Hydrogen polonide

 Hydrogen peroxide

Related  Acetone
solvents
 Ethanol

 Methanol

 Hydrogen fluoride

 Ammonia

Supplementary data page

Water (data page)

Except where otherwise noted, data are given for

materials in their standard state (at 25 °C [77 °F],
100 kPa).

verify (what is ?)

Infobox references

Water is an inorganic compound with the chemical formula H2O. It is a transparent,

tasteless, odorless,[c] and nearly colorless chemical substance. It is the main constituent
of Earth's hydrosphere and the fluids of all known living organisms (in which it acts as
a solvent[20]). It is vital for all known forms of life, despite not providing food energy or
organic micronutrients. Its chemical formula, H2O, indicates that each of
its molecules contains one oxygen and two hydrogen atoms, connected by covalent bonds.
The hydrogen atoms are attached to the oxygen atom at an angle of 104.45°. [21] In liquid
form, H2O is also called "water" at standard temperature and pressure.

Because Earth's environment is relatively close to water's triple point, water exists on Earth
as a solid, a liquid, and a gas.[22] It forms precipitation in the form of rain and aerosols in the
form of fog. Clouds consist of suspended droplets of water and ice, its solid state. When
finely divided, crystalline ice may precipitate in the form of snow. The gaseous state of water
is steam or water vapor.

Water covers about 71% of the Earth's surface, with seas and oceans making up most of the
water volume (about 96.5%).[23] Small portions of water occur as groundwater (1.7%), in
the glaciers and the ice caps of Antarctica and Greenland (1.7%), and in the air as vapor,
clouds (consisting of ice and liquid water suspended in air), and precipitation (0.001%). [24]
[25]
Water moves continually through the water
cycle of evaporation, transpiration (evapotranspiration), condensation, precipitation,
and runoff, usually reaching the sea.

Water plays an important role in the world economy. Approximately 70% of the fresh
water used by humans goes to agriculture.[26] Fishing in salt and fresh water bodies has been,
and continues to be, a major source of food for many parts of the world, providing 6.5% of
global protein.[27] Much of the long-distance trade of commodities (such as oil, natural gas,
and manufactured products) is transported by boats through seas, rivers, lakes, and canals.
Large quantities of water, ice, and steam are used for cooling and heating in industry and
homes. Water is an excellent solvent for a wide variety of substances, both mineral and
organic; as such, it is widely used in industrial processes and in cooking and washing. Water,
ice, and snow are also central to many sports and other forms of entertainment, such
as swimming, pleasure boating, boat racing, surfing, sport fishing, diving, ice
skating, snowboarding, and skiing.

Etymology

The word water comes from Old English wæter, from Proto-Germanic *watar (source also
of Old Saxon watar, Old Frisian wetir, Dutch water, Old High
German wazzar, German Wasser, vatn, Gothic 𐍅𐌰𐍄𐍉 (wato)), from Proto-Indo-
European *wod-or, suffixed form of root *wed- ('water'; 'wet').[28] Also cognate, through the
Indo-European root, with Greek ύδωρ (ýdor; from Ancient Greek ὕδωρ (hýdōr), whence
English 'hydro-'), Russian вода́ (vodá), Irish uisce, and Albanian ujë.

History

Main articles: Origin of water on Earth § History of water on Earth, and Properties of water
§ History

On Earth

This section is an excerpt from Origin of water on Earth § History of water on Earth.[edit]

One factor in estimating when water appeared on Earth is that water is continually being lost
to space. H2O molecules in the atmosphere are broken up by photolysis, and the resulting
free hydrogen atoms can sometimes escape Earth's gravitational pull. When the Earth was
younger and less massive, water would have been lost to space more easily. Lighter
elements like hydrogen and helium are expected to leak from the atmosphere continually,
but isotopic ratios of heavier noble gases in the modern atmosphere suggest that even the
heavier elements in the early atmosphere were subject to significant losses.[29] In
particular, xenon is useful for calculations of water loss over time. Not only is it a noble gas
(and therefore is not removed from the atmosphere through chemical reactions with other
elements), but comparisons between the abundances of its nine stable isotopes in the
modern atmosphere reveal that the Earth lost at least one ocean of water early in its history,
between the Hadean and Archean eons.[30][clarification needed]

Any water on Earth during the latter part of its accretion would have been disrupted by
the Moon-forming impact (~4.5 billion years ago), which likely vaporized much of Earth's
crust and upper mantle and created a rock-vapor atmosphere around the young planet.[31]
[32]
The rock vapor would have condensed within two thousand years, leaving behind hot
volatiles which probably resulted in a majority carbon dioxide atmosphere with hydrogen
and water vapor. Afterward, liquid water oceans may have existed despite the surface
temperature of 230 °C (446 °F) due to the increased atmospheric pressure of the
CO2 atmosphere. As the cooling continued, most CO2 was removed from the atmosphere
by subduction and dissolution in ocean water, but levels oscillated wildly as new surface
and mantle cycles appeared.[33]

This pillow basalt on the seafloor near Hawaii was formed

when magma extruded underwater. Other, much older pillow basalt formations provide
evidence for large bodies of water long ago in Earth's history.

Geological evidence also helps constrain the time frame for liquid water existing on Earth. A
sample of pillow basalt (a type of rock formed during an underwater eruption) was
recovered from the Isua Greenstone Belt and provides evidence that water existed on Earth
3.8 billion years ago.[34] In the Nuvvuagittuq Greenstone Belt, Quebec, Canada, rocks dated
at 3.8 billion years old by one study[35] and 4.28 billion years old by another[36] show evidence
of the presence of water at these ages.[34] If oceans existed earlier than this, any geological
evidence has yet to be discovered (which may be because such potential evidence has been
destroyed by geological processes like crustal recycling). More recently, in August 2020,
researchers reported that sufficient water to fill the oceans may have always been on
the Earth since the beginning of the planet's formation.[37][38][39]

Unlike rocks, minerals called zircons are highly resistant to weathering and geological
processes and so are used to understand conditions on the very early Earth. Mineralogical
evidence from zircons has shown that liquid water and an atmosphere must have existed
4.404 ± 0.008 billion years ago, very soon after the formation of Earth.[40][41][42][43] This
presents somewhat of a paradox, as the cool early Earth hypothesis suggests temperatures
were cold enough to freeze water between about 4.4 billion and 4.0 billion years ago. Other
studies of zircons found in Australian Hadean rock point to the existence of plate tectonics as
early as 4 billion years ago. If true, that implies that rather than a hot, molten surface and an
atmosphere full of carbon dioxide, early Earth's surface was much as it is today (in terms
of thermal insulation). The action of plate tectonics traps vast amounts of CO2, thereby
reducing greenhouse effects, leading to a much cooler surface temperature and the
formation of solid rock and liquid water.[44]

Properties

Main article: Properties of water

See also: Water (data page) and Water model

A water molecule consists of two hydrogen atoms and one

oxygen atom.

Water (H2O) is a polar inorganic compound. At room temperature it is

a tasteless and odorless liquid, nearly colorless with a hint of blue. The simplest hydrogen
chalcogenide, it is by far the most studied chemical compound and is sometimes described
as the "universal solvent" for its ability to dissolve more substances than any other liquid,[45]
[46]
though it is poor at dissolving nonpolar substances.[47] This allows it to be the "solvent of
life":[48] indeed, water as found in nature almost always includes various dissolved
substances, and special steps are required to obtain chemically pure water. Water is the only
common substance to exist as a solid, liquid, and gas in normal terrestrial conditions.[49]

States

The three common states of matter

Along with oxidane, water is one of the two official names for the chemical compound H
[50]
2O; it is also the liquid phase of H
[51]
2O. The other two common states of matter of water are the solid phase, ice, and the
gaseous phase, water vapor or steam. The addition or removal of heat can cause phase
transitions: freezing (water to ice), melting (ice to water), vaporization (water to
vapor), condensation (vapor to water), sublimation (ice to vapor) and deposition (vapor to
ice).[52]

Density

See also: Frost weathering

Water is one of only a few common naturally occurring substances which, for some
temperature ranges, become less dense as they cool, and the only known naturally occurring
substance which does so while liquid. In addition it is unusual as it becomes significantly
less dense as it freezes, though it is not unique in that respect.[d]

At 1 atm pressure, it reaches its maximum density of 999.972 kg/m3 (62.4262 lb/cu ft) at
3.98 °C (39.16 °F).[54][55]

Below that temperature, but above the freezing point of 0 °C (32 °F), it expands becoming
less dense until it reaches freezing point, reaching a density in its liquid phase of
999.8 kg/m3 (62.4155 lb/cu ft).

Once it freezes and becomes ice, it expands by about 9%, with a density of
917 kg/m3 (57.25 lb/cu ft).[56][57] This expansion can exert enormous pressure, bursting pipes
and cracking rocks.[58] As a solid, it displays the usual behavior of contracting and becoming
more dense as it cools. These unusual thermal properties have important consequences for
life on earth.

In a lake or ocean, water at 4 °C (39 °F) sinks to the bottom, and ice forms on the surface,
floating on the liquid water. This ice insulates the water below, preventing it from freezing
solid. Without this protection, most aquatic organisms residing in lakes would perish during
the winter.[59] In addition, this anomalous behavior is an important part of the thermohaline
circulation which distributes heat around the planet's oceans.

Magnetism

Water is a diamagnetic material.[60] Though interaction is weak, with superconducting

magnets it can attain a notable interaction.[60]

Phase transitions

At a pressure of one atmosphere (atm), ice melts or water freezes (solidifies) at 0 °C (32 °F)
and water boils or vapor condenses at 100 °C (212 °F). However, even below the boiling
point, water can change to vapor at its surface by evaporation (vaporization throughout the
liquid is known as boiling). Sublimation and deposition also occur on surfaces.[52] For
example, frost is deposited on cold surfaces while snowflakes form by deposition on an
aerosol particle or ice nucleus.[61] In the process of freeze-drying, a food is frozen and then
stored at low pressure so the ice on its surface sublimates.[62]

The melting and boiling points depend on pressure. A good approximation for the rate of
change of the melting temperature with pressure is given by the Clausius–Clapeyron
relation:

where and are the molar volumes of the liquid and solid phases, and is the molar latent
heat of melting. In most substances, the volume increases when melting occurs, so the
melting temperature increases with pressure. However, because ice is less dense than water,
the melting temperature decreases.[53] In glaciers, pressure melting can occur under
sufficiently thick volumes of ice, resulting in subglacial lakes.[63][64]

The Clausius-Clapeyron relation also applies to the boiling point, but with the liquid/gas
transition the vapor phase has a much lower density than the liquid phase, so the boiling
point increases with pressure.[65] Water can remain in a liquid state at high temperatures in
the deep ocean or underground. For example, temperatures exceed 205 °C (401 °F) in Old
Faithful, a geyser in Yellowstone National Park.[66] In hydrothermal vents, the temperature
can exceed 400 °C (752 °F).[67]

At sea level, the boiling point of water is 100 °C (212 °F). As atmospheric pressure decreases
with altitude, the boiling point decreases by 1 °C every 274 meters. High-altitude
cooking takes longer than sea-level cooking. For example, at 1,524 metres (5,000 ft), cooking
time must be increased by a fourth to achieve the desired result.[68] Conversely, a pressure
cooker can be used to decrease cooking times by raising the boiling temperature.[69] In a
vacuum, water will boil at room temperature.[70]

Triple and critical points

Phase diagram of water

On a pressure/temperature phase diagram (see figure), there are curves separating solid
from vapor, vapor from liquid, and liquid from solid. These meet at a single point called
the triple point, where all three phases can coexist. The triple point is at a temperature of
273.16 K (0.01 °C; 32.02 °F) and a pressure of 611.657 pascals (0.00604 atm; 0.0887 psi);[71] it
is the lowest pressure at which liquid water can exist. Until 2019, the triple point was used to
define the Kelvin temperature scale.[72][73]

The water/vapor phase curve terminates at 647.096 K (373.946 °C; 705.103 °F) and 22.064
megapascals (3,200.1 psi; 217.75 atm).[74] This is known as the critical point. At higher
temperatures and pressures the liquid and vapor phases form a continuous phase called
a supercritical fluid. It can be gradually compressed or expanded between gas-like and
liquid-like densities; its properties (which are quite different from those of ambient water)
are sensitive to density. For example, for suitable pressures and temperatures it can mix
freely with nonpolar compounds, including most organic compounds. This makes it useful in
a variety of applications including high-temperature electrochemistry and as an ecologically
benign solvent or catalyst in chemical reactions involving organic compounds. In Earth's
mantle, it acts as a solvent during mineral formation, dissolution and deposition. [75][76]

Phases of ice and water

Main article: Ice

The normal form of ice on the surface of Earth is ice Ih, a phase that forms crystals
with hexagonal symmetry. Another with cubic crystalline symmetry, ice Ic, can occur in the
upper atmosphere.[77] As the pressure increases, ice forms other crystal structures. As of
2024, twenty have been experimentally confirmed and several more are predicted
theoretically.[78] The eighteenth form of ice, ice XVIII, a face-centred-cubic, superionic ice
phase, was discovered when a droplet of water was subject to a shock wave that raised the
water's pressure to millions of atmospheres and its temperature to thousands of degrees,
resulting in a structure of rigid oxygen atoms in which hydrogen atoms flowed freely.[79]
[80]
When sandwiched between layers of graphene, ice forms a square lattice.[81]

The details of the chemical nature of liquid water are not well understood; some theories
suggest that its unusual behavior is due to the existence of two liquid states.[55][82][83][84]

Taste and odor

Pure water is usually described as tasteless and odorless, although humans have specific
sensors that can feel the presence of water in their mouths,[85][86] and frogs are known to be
able to smell it.[87] However, water from ordinary sources (including mineral water) usually
has many dissolved substances that may give it varying tastes and odors. Humans and other
animals have developed senses that enable them to evaluate the potability of water in order
to avoid water that is too salty or putrid.[88]

Color and appearance

Main article: Color of water

See also: Electromagnetic absorption by water

Pure water is visibly blue due to absorption of light in the region c. 600–800 nm.[89] The color
can be easily observed in a glass of tap-water placed against a pure white background, in
daylight. The principal absorption bands responsible for the color are overtones of the O–H
stretching vibrations. The apparent intensity of the color increases with the depth of the
water column, following Beer's law. This also applies, for example, with a swimming pool
when the light source is sunlight reflected from the pool's white tiles.

In nature, the color may also be modified from blue to green due to the presence of
suspended solids or algae.

In industry, near-infrared spectroscopy is used with aqueous solutions as the greater

intensity of the lower overtones of water means that glass cuvettes with short path-length
may be employed. To observe the fundamental stretching absorption spectrum of water or
of an aqueous solution in the region around 3,500 cm−1 (2.85 μm)[90] a path length of about
25 μm is needed. Also, the cuvette must be both transparent around 3500 cm−1 and
insoluble in water; calcium fluoride is one material that is in common use for the cuvette
windows with aqueous solutions.

The Raman-active fundamental vibrations may be observed with, for example, a 1 cm

sample cell.

Aquatic plants, algae, and other photosynthetic organisms can live in water up to hundreds
of meters deep, because sunlight can reach them. Practically no sunlight reaches the parts of
the oceans below 1,000 metres (3,300 ft) of depth.

The refractive index of liquid water (1.333 at 20 °C (68 °F)) is much higher than that of air
(1.0), similar to those of alkanes and ethanol, but lower than those
of glycerol (1.473), benzene (1.501), carbon disulfide (1.627), and common types of glass
(1.4 to 1.6). The refraction index of ice (1.31) is lower than that of liquid water.

Molecular polarity

Tetrahedral structure of water

In a water molecule, the hydrogen atoms form a 104.5° angle with the oxygen atom. The
hydrogen atoms are close to two corners of a tetrahedron centered on the oxygen. At the
other two corners are lone pairs of valence electrons that do not participate in the bonding.
In a perfect tetrahedron, the atoms would form a 109.5° angle, but the repulsion between
the lone pairs is greater than the repulsion between the hydrogen atoms.[91][92] The O–H
bond length is about 0.096 nm.[93]

Other substances have a tetrahedral molecular structure, for example methane (CH
4) and hydrogen sulfide (H

2S). However, oxygen is more electronegative than most other elements, so the oxygen atom

has a negative partial charge while the hydrogen atoms are partially positively charged.
Along with the bent structure, this gives the molecule an electrical dipole moment and it is
classified as a polar molecule.[94]

Water is a good polar solvent, dissolving many salts and hydrophilic organic molecules such
as sugars and simple alcohols such as ethanol. Water also dissolves many gases, such as
oxygen and carbon dioxide—the latter giving the fizz of carbonated beverages, sparkling
wines and beers. In addition, many substances in living organisms, such
as proteins, DNA and polysaccharides, are dissolved in water. The interactions between
water and the subunits of these biomacromolecules shape protein folding, DNA base
pairing, and other phenomena crucial to life (hydrophobic effect).

Many organic substances (such as fats and oils and alkanes) are hydrophobic, that is,
insoluble in water. Many inorganic substances are insoluble too, including most
metal oxides, sulfides, and silicates.

Hydrogen bonding

See also: Chemical bonding of water

Model of hydrogen bonds (1) between molecules of water

Because of its polarity, a molecule of water in the liquid or solid state can form up to
four hydrogen bonds with neighboring molecules. Hydrogen bonds are about ten times as
strong as the Van der Waals force that attracts molecules to each other in most liquids. This
is the reason why the melting and boiling points of water are much higher than those
of other analogous compounds like hydrogen sulfide. They also explain its exceptionally
high specific heat capacity (about 4.2 J/(g·K)), heat of fusion (about 333 J/g), heat of
vaporization (2257 J/g), and thermal conductivity (between 0.561 and 0.679 W/(m·K)).
These properties make water more effective at moderating Earth's climate, by storing heat
and transporting it between the oceans and the atmosphere. The hydrogen bonds of water
are around 23 kJ/mol (compared to a covalent O-H bond at 492 kJ/mol). Of this, it is
estimated that 90% is attributable to electrostatics, while the remaining 10% is partially
covalent.[95]

These bonds are the cause of water's high surface tension[96] and capillary forces.
The capillary action refers to the tendency of water to move up a narrow tube against the
force of gravity. This property is relied upon by all vascular plants, such as trees.[citation needed]