| Water (H2O) |
|---|
 The water molecule has this basic geometric structure  Water molecules have this structure. |
| General | |
|---|---|
| Systematic name | Water |
| Other names | Aqua
Hydrogen oxide
Hydrogen hydroxide
Hydrate
Oxidane
Hydric acid
Dihydrogen monoxide
Hydroxyl acid
Dihydrogen oxide
Hydrohydroxic acid
''μ''-Oxido dihydrogen
Light Water |
| Molecular formula | HOH or H2O |
| Molar mass | 18.01524 g·mol−1 |
| Appearance | transparent, almost
colorless liquid with
a slight hint of blue[1] |
| CAS number | [7732-18-5] |
| see also | 'Water (data page)' |
| Properties |
|---|
| Density and phase | 1000 kg·m−3, liquid (4 °C)
917 kg·m−3, solid |
| Melting point | 0 °C, 32 °F (273.15 K)[2] |
| Boiling point | 100 °C, 212 °F (373.15 K) |
| Triple point | 273.16 K, 611.73 Pa |
| Critical point | 647 K, 22.1 MPa |
Specific heat
capacity ''(gas)'' | ''cp''=1970 J·kg−1·K−1 @ 300 K
''cv''=1510 J·kg−1·K−1 @ 300 K[3] |
Specific heat
capacity ''(liquid)'' | 4186 J·kg−1·K−1 |
Specific heat
capacity ''(solid)'' | 2060 J·kg−1·K−1 |
| Acidity (p''K''a) | 15.74
~22 |
| Basicity (p''K''b) | 15.74 |
| Viscosity | 0.001 Pa·s at 20 °C |
| Surface Tension at 20 °C | 72.86 mN·m−1 |
| Structure |
|---|
| Molecular shape | non-linear bent |
| Point Group | C2v |
| Crystal structure | Hexagonal
''See ice'' |
| Dipole moment | 1.85 D |
| Hazards |
|---|
| MSDS | External MSDS |
| Main hazards | see Dihydrogen monoxide hoax |
| NFPA 704 | |
| RTECS number | ZC0110000 |
| Supplementary data page |
|---|
Structure and
properties | ''n'', ''εr'', etc. |
Thermodynamic
data | Phase behaviour
Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| Related compounds |
|---|
| Related solvents | acetone
methanol |
| Related compounds | water vapor
ice
heavy water |
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
|
'Water' (
H2O, HOH) is the most abundant molecule on Earth's surface, composing of about 70% of the Earth's surface as liquid and solid state in addition to being found in the atmosphere as a vapor. It is in
dynamic equilibrium between the
liquid and
vapor states at
standard temperature and pressure. At
room temperature, it is a nearly
colorless,
tasteless, and
odorless liquid. Many substances dissolve in water and it is commonly referred to as ''the universal
solvent''. Because of this, water in nature and in use is rarely clean, and may have some properties different than those in the laboratory. However, there are many compounds that are essentially, if not completely, insoluble in water. Water is the only common, pure substance found naturally in all three
states of matter—for other substances, see
Chemical properties.
Forms of water
:''See the ''
Water can take many forms. The
solid state of water is commonly known as
ice (while many other forms exist; see
amorphous solid water); the
gaseous state is known as
water vapor (or
steam, though this is actually incorrect, since steam is just condensing ''liquid'' water droplets), and the common liquid
phase is generally taken as simply water. Above a certain
critical temperature and pressure (647
K and 22.064
MPa), water molecules assume a ''supercritical'' condition, in which liquid-like clusters float within a vapor-like phase.
Heavy water is water in which the hydrogen is replaced by its heavier
isotope,
deuterium. It is ''chemically'' almost identical to normal water. Heavy water is used in the
nuclear industry to slow down
neutrons.
Water in the Universe
Water has been detected in
interstellar clouds within our
galaxy, the
Milky Way. It is believed that water exists in abundance in other galaxies too, because its components,
hydrogen and
oxygen, are among the most abundant elements in the universe.
Interstellar clouds eventually condense into
solar nebulae and
solar systems, such as ours. The initial water can then be found in
comets,
planets,
dwarf planets, and their
satellites. In our solar system, water, in ice form, has been found:
★ on the
Moon,
★ on the planets
Mercury,
Mars, and
Neptune,
★ on the dwarf planet
Pluto,
★ on satellites of planets, such as
Triton and
Europa.
The liquid form of water is only known to occur on Earth, though strong evidence suggests that it is present just under the surface of
Saturn's moon
Enceladus.
Water on Earth
The
water cycle (known scientifically as the 'hydrologic cycle') refers to the continuous exchange of water within the
hydrosphere, between the
atmosphere,
soil water,
surface water,
groundwater, and
plants.
Earth's approximate water volume (the total water supply of the world) is 1 360 000 000 km
3 (326 000 000 mi
3). Of this volume:
★ 1 320 000 000 km
3 (316 900 000 mi
3 or 97.2%) is in the
oceans.
★ 25 000 000 km
3 (6 000 000 mi
3 or 1.8%) is in
glaciers,
ice caps and
ice sheets.
★ 13 000 000 km
3 (3,000,000 mi
3 or 0.9%) is
groundwater.
★ 250 000 km
3 (60,000 mi
3 or 0.02%) is
fresh water in lakes, inland seas, and rivers.
★ 13 000 km
3 (3,100 mi
3 or 0.001%) is atmospheric water vapor at any given time.
Liquid water is found in 'bodies of water', such as an
ocean,
sea,
lake,
river,
stream,
canal,
pond, or
puddle. The majority of water on Earth is
sea water. Water is also present in the atmosphere in solid, liquid, and vapor phases. It also exists as groundwater in
aquifers.
The boiling point of water is directly related to the barometric pressure. For example, on the top of
Mt. Everest water boils at about 68 degrees Celsius, compared to 100 degrees at
sea level. Conversely, water deep in the ocean near geothermal vents can reach temperatures of hundreds of degrees and remain liquid.
Water in industry
Water is also used in many industrial processes and machines, such as the
steam turbine and
heat exchanger, in addition to its use as a chemical
solvent. Discharge of untreated water from industrial uses is
pollution. Pollution includes discharged solutes (
chemical pollution) and discharged coolant water (thermal pollution). Industry requires pure water for many applications and utilizes a variety of
purification techniques both in water supply and discharge.
Physics and chemistry of water
Density of water and ice
The solid form of most substances is more
dense than the liquid
phase; thus, a block of pure solid ''substance'' will sink in a tub of pure liquid ''substance''. But, by contrast, a block of common
ice will float in a tub of water because solid water is ''less'' dense than liquid water. This is an extremely important characteristic property of water. At
room temperature, liquid water becomes denser with lowering temperature, just like other substances. But at 4 °C, just above freezing, water reaches its
maximum density, and as water cools further toward its freezing point, the liquid water, under standard conditions, expands to become ''less'' dense. The physical reason for this is related to the
crystal structure of ordinary
ice, known as
hexagonal ice Ih. Water,
gallium,
bismuth,
antimony and
silicon are some of the few materials which expand when they freeze; most other materials contract. It should be noted however, that not all forms of ice are less dense than liquid water. For example
HDA and
VHDA are both more dense than liquid phase pure water. Thus, the reason that the common form of ice is less dense than water is a bit non-intuitive and relies heavily on the unusual properties inherent to the
hydrogen bond.
Generally, water expands when it freezes because of its
molecular structure, in tandem with the unusual
elasticity of the hydrogen bond and the particular lowest energy hexagonal
crystal conformation that it adopts under standard conditions. That is, when water cools, it tries to stack in a
crystalline lattice configuration that stretches the
rotational and
vibrational components of the bond, so that the effect is that each molecule of water is pushed further from each of its neighboring molecules. This effectively reduces the density ''ρ'' of water when ice is formed under standard conditions.
The importance of this property cannot be overemphasized for its role on the
ecosystem of Earth. For example, ''if'' water were more dense when frozen, lakes and oceans in a polar environment would eventually freeze solid (from top to bottom). This would happen because frozen ice would settle on the lake and riverbeds, and the necessary warming phenomenon (see below) could not occur in summer, as the warm surface layer would be less dense than the solid frozen layer below. It is a significant feature of nature that this does not occur naturally in the environment.
Nevertheless, the unusual expansion of freezing water (in ordinary ''natural'' settings in relevant biological systems), due to the
hydrogen bond, from 4 °C above freezing to the freezing point offers an important advantage for freshwater life in winter. Water chilled at the surface increases in density and sinks, forming
convection currents that cool the whole water body, but when the temperature of the lake water reaches 4 °C, water on the surface decreases in density as it chills further and remains as a surface layer which eventually
freezes and forms ice. Since downward convection of colder water is blocked by the density change, any large body of fresh water frozen in winter will have the coldest water near the surface, away from the
riverbed or lakebed. This accounts for various little known phenomena of ice characteristics as they relate to ice in lakes and "ice falling out of lakes" as described by early 20th century scientist Horatio D. Craft.
The following table gives the density of water in grams per cubic centimeter at various temperatures in degrees Celsius:
[4]
| Temp (°C) | Density (g/cm3) |
|---|
| 30 | 0.9956473 |
| 20 | 0.9982041 |
| 10 | 0.9996996 |
| 4 | 0.9999720 |
| 0 | 0.9998395 |
| −10 | 0.998117 |
| −20 | 0.993547 |
| −30 | 0.983854 |
The values below 0 °C refer to
supercooled water.
Density of saltwater and ice
The density of water is dependent on the temperature of the water. This is because the density is different for salt water than for fresh water. Ice still floats in the oceans, otherwise they would freeze from the bottom up. However, the salt content of oceans lowers the freezing point by about 2 °C and lowers the temperature of the density maximum of water to the freezing point. That is why, in ocean water, the downward convection of colder water is ''not'' blocked by an expansion of water as it becomes colder near the freezing point. The oceans' cold water near the freezing point continues to sink. For this reason, any creature attempting to survive at the bottom of such cold water as the
Arctic Ocean generally lives in water that is 4 °C colder than the temperature at the bottom of frozen-over
fresh water lakes and rivers in winter.
As the
surface of salt water begins to freeze (at −1.9 °C for normal salinity
seawater, 3.5%) the ice that forms is essentially salt free with a density approximately equal to that of freshwater ice. This ice floats on the surface and the salt that is "frozen out" adds to the
salinity and density of the seawater just below it, in a process known as ''
brine rejection''. This more dense saltwater sinks by convection and the replacing seawater is subject to the same process. This provides essentially freshwater ice at −1.9 °C on the surface. The increased density of the seawater beneath the forming ice causes it to sink towards the bottom.
Compressibility
The
compressibility of water is a function of pressure and temperature. At 0 °C in the limit of zero pressure the compressibility is 5.1×10
7 bar
−1.
[5] In the zero pressure limit the compressibility reaches a minimum of 4.4×10
7 bar
−1 around 45 °C before increasing again with increasing temperature. As the pressure is increased the compressibility decreases, being 3.9×10
7 bar
−1 at 0 °C and 1000
bar.
The
bulk modulus of water is 2.2×10
9 Pa.
[6] The low compressibility of non-gases, and of water in particular, leads to them often being incorrectly labelled as incompressible. The low compressibility of water means that even in the deep
oceans at 4000
m depth, where pressures are 4×10
7 Pa, there is only a 1.8% decrease in volume.
[6]
Triple point
The
temperature and
pressure at which solid, liquid, and
gaseous water coexist in equilibrium is called the
triple point of water. This point is used to define the units of temperature (the
kelvin and, indirectly, the degree
Celsius and even the degree
Fahrenheit). The triple point is at a temperature of 273.16 K (0.01 °C) by convention, and at a pressure of 611.73
Pa. This pressure is quite low, about 1/166 of the normal sea level barometric pressure of 101,325 Pa. The atmospheric surface pressure on planet
Mars is remarkably close to the triple point pressure, and the zero-elevation or "sea level" of
Mars is defined by the height at which the atmospheric pressure corresponds to the triple point of water.
Mpemba effect
The
Mpemba effect is the surprising phenomenon whereby hot water can, under certain conditions, freeze sooner than cold water, even though it must pass the lower temperature on the way to freezing. However, this can be explained with
evaporation,
convection,
supercooling, and the
insulating effect of
frost.
Hot ice
Hot ice is the name given to another surprising phenomenon in which water at room temperature can be turned into ice ''that remains at room temperature'' by supplying an electric field on the order of 10
6 volts per meter.
[8]
The effect of such electric fields has been suggested as an explanation of cloud formation. The first time cloud ice forms around a clay particle, it requires a temperature of −10 °C, but subsequent freezing around the same clay particle requires a temperature of just −5 °C, suggesting some kind of structural change.
[9]
Surface tension
Water drops are stable, due to the high
surface tension of water, 72.8 mN/m, the highest of the non-metallic liquids. This can be seen when small quantities of water are put on a surface such as glass: the water stays together as drops. This property is important for life. For example, when water is carried through
xylem up stems in plants the strong intermolecular attractions hold the water column together. Strong cohesive properties hold the water column together, and strong adhesive properties stick the water to the xylem, and prevent tension rupture caused by
transpiration pull. Other liquids with lower surface tension would have a higher tendency to "rip", forming vacuum or air pockets and rendering the xylem water transport inoperative.
Electrical properties
Pure water containing no ions is an excellent
insulator, however, not even "deionized" water, is completely free of ions. Water undergoes
auto-ionisation at any temperature above
absolute zero. Further, because water is such a good solvent, it almost always has some
solute dissolved in it, most frequently a
salt. If water has even a tiny amount of such an impurity, then it can conduct electricity readily, as impurities such as salt separate into free
ions in aqueous solution by which an electric current can flow.
Water can be split into its constituent elements, hydrogen and oxygen, by passing a current through it. This process is called
electrolysis. Water molecules naturally dissociate into H
+ and OH
− ions, which are pulled toward the
cathode and
anode, respectively. At the cathode, two H
+ ions pick up electrons and form H
2 gas. At the anode, four OH
− ions combine and release O
2 gas, molecular water, and four electrons. The gases produced bubble to the surface, where they can be collected. It is known that the theoretical maximum electrical resistivity for water is approximately 182
kilohm-m
2/m (or 18.2 MΩ·cm
2/cm) at 25 °C. This figure agrees well with what is typically seen on
reverse osmosis, ultrafiltered and deionized
ultrapure water systems used for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding that of even 100 parts per trillion (ppt) in ultrapure water will begin to noticeably lower its resistivity level by up to several kilohm-square meters/meter (a change of several hundred
nanosiemens per meter of conductance).
Dipolar nature of water
An important feature of water is its
polar nature. The water molecule forms an angle, with hydrogen atoms at the tips and oxygen at the vertex. Since oxygen has a higher
electronegativity than hydrogen, the side of the molecule with the oxygen atom has a partial negative charge. A molecule with such a charge difference is called a
dipole. The charge differences cause water molecules to be attracted to each other (the relatively positive areas being attracted to the relatively negative areas) and to other polar molecules. This attraction is known as
hydrogen bonding, and explains many of the properties of water. Certain molecules, such as carbon dioxide, also have a difference in electronegativity between the atoms but the difference is that the shape of carbon dioxide is symmetrically aligned and so the opposing charges cancel one another out. This phenomenon of water can be seen if you hold an electrical source near a thin stream of water falling vertically, causing the stream to bend towards the electrical source.
Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for a number of water's physical properties. One such property is its relatively high
melting and
boiling point temperatures; more
heat energy is required to break the hydrogen bonds between molecules. The similar compound hydrogen sulfide (H
2S), which has much weaker hydrogen bonding, is a gas at
room temperature even though it has twice the molecular mass of water. The extra bonding between water molecules also gives liquid water a large
specific heat capacity. This high heat capacity makes water a good heat storage medium.
Hydrogen bonding also gives water its unusual behavior when freezing. When cooled to near freezing point, the presence of hydrogen bonds means that the molecules, as they rearrange to minimize their energy, form the
hexagonal crystal structure of ice that is actually of lower density: hence the solid form, ice, will float in water. In other words, water expands as it freezes, whereas almost all other materials shrink on solidification.
An interesting consequence of the solid having a lower density than the liquid is that ice will melt if sufficient pressure is applied. With increasing pressure the melting point temperature drops and when the melting point temperature is lower than the ambient temperature the ice begins to melt. A significant increase of pressure is required to lower the melting point temperature by very much—the pressure exerted by an ice skater on the ice would only reduce the melting point by approximately 0.09 °C.
Water as a solvent
Water is also a good
solvent due to its
polarity. When an ionic or polar compound enters water, it is surrounded by water molecules (
Hydration). The relatively small size of water molecules typically allows many water molecules to surround one molecule of
solute. The partially negative dipole ends of the water are attracted to positively charged components of the solute, and vice versa for the positive dipole ends.
In general, ionic and polar substances such as
acids,
alcohols, and
salts are relatively soluble in water, and nonpolar substances such as fats and oils are not. Nonpolar molecules stay together in water because it is energetically more favorable for the water molecules to hydrogen bond to each other than to engage in
van der Waals interactions with nonpolar molecules.
An example of an ionic solute is
table salt; the sodium chloride, NaCl, separates into Na
+ cations and Cl
- anions, each being surrounded by water molecules. The ions are then easily transported away from their
crystalline lattice into solution. An example of a nonionic solute is
table sugar. The water dipoles make hydrogen bonds with the polar regions of the sugar molecule (OH groups) and allow it to be carried away into solution.
Amphoteric nature of water
Chemically, water is
amphoteric — i.e., it is able to act as either an
acid or a
base. Occasionally the term ''hydroxic acid'' is used when water acts as an acid in a chemical reaction. At a pH of 7 (neutral), the concentration of
hydroxide ions (OH
−) is equal to that of the
hydronium (H
3O
+) or
hydrogen (H
+) ions. If the
equilibrium is disturbed, the solution becomes acidic (higher concentration of hydronium ions) or basic (higher concentration of hydroxide ions).
Water can act as either an acid or a base in reactions. According to the
Brønsted-Lowry system, an acid is defined as a species which donates a proton (an H
+ ion) in a reaction, and a base as one which receives a proton. When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid. For instance, it receives an H
+ ion from HCl in the equilibrium:
:HCl + H
2O H
3O
+ + Cl
−
Here water is acting as a base, by receiving an H
+ ion.
In the reaction with
ammonia, NH
3, water donates an H
+ ion, and is thus acting as an acid:
:NH
3 + H
2O NH
4+ + OH
−
Acidity in nature
In theory, pure water has a
pH of 7 at 298 K. In practice, pure water is very difficult to produce. Water left exposed to air for any length of time will rapidly dissolve
carbon dioxide, forming a dilute solution of
carbonic acid, with a limiting pH of about 5.7. As cloud droplets form in the atmosphere and as raindrops fall through the air minor amounts of CO
2 are absorbed and thus most rain is slightly acidic. If high amounts of
nitrogen and
sulfur oxides are present in the air, they too will dissolve into the cloud and rain drops producing more serious
acid rain problems.
Hydrogen bonding in water
A water molecule can form a maximum of four
hydrogen bonds because it can accept two and donate two hydrogens. Other molecules like
hydrogen fluoride,
ammonia,
methanol form hydrogen bonds but they do not show anomalous behaviour of
thermodynamic,
kinetic or structural properties like those observed in water. The answer to the apparent difference between water and other hydrogen bonding liquids lies in the fact that apart from water none of the hydrogen bonding molecules can form four hydrogen bonds either due to an inability to donate/accept hydrogens or due to
steric effects in bulky residues. In water local
tetrahedral order due to the four hydrogen bonds gives rise to an open structure and a 3-dimensional bonding network, which exists in contrast to the closely packed structures of simple
liquids. There is a great similarity between water and
silica in their anomalous behaviour, even though one (water) is a liquid which has a hydrogen bonding network while the other (silica) has a covalent network with a very high melting point. One reason that water is well suited, and chosen, by life-forms, is that it exhibits its unique properties over a temperature regime that suits diverse
biological processes, including
hydration.
It is believed that hydrogen bond in water is largely due to electrostatic forces and some amount of covalency. The partial covalent nature of hydrogen bond predicted by
Linus Pauling in 1930s is yet to be proven unambiguously by experiments and theoretical calculations.
Quantum properties of molecular water
Although the molecular formula of water is generally considered to be a stable result in molecular thermodynamics, recent work, started in 1995 has shown that at certain scales, water may act more like H
3/2O than H
2O at the subatomic quantum level.
[10] This result could have significant ramifications at the level of, for example, the
hydrogen bond in
biological,
chemical and
physical systems. The experiment shows that when
neutrons and
protons collide with water, they scatter in a way that indicates that they only are affected by a ratio of 1.5:1 of
hydrogen to
oxygen respectively. However, the time-scale of this response is only seen at the level of attoseconds (10
-18 seconds), and so is only relevant in highly resolved
kinetic and
dynamical systems.
[11][12]
History
The properties of water have historically been used to define various
temperature scales. Notably, the
Kelvin,
Celsius and
Fahrenheit scales were, or currently are, defined by the freezing and boiling points of water. The less common scales of
Delisle,
Newton,
Réaumur and
Rømer were defined similarly. The
triple point of water is a more commonly used standard point today.
[13]
The first scientific decomposition of water into hydrogen and oxygen, by
electrolysis, was done in 1800 by
William Nicholson, an English chemist. In 1805,
Joseph Louis Gay-Lussac and
Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen (by volume).
Gilbert Newton Lewis isolated the first sample of pure
heavy water in 1933.
Polywater was a hypothetical
polymerized form of water that was the subject of much scientific controversy during the late 1960s. The consensus now is that it does not exist.
Systematic naming
The accepted
IUPAC name of water is simply "water", although there are two other systematic names which can be used to describe the molecule.
The simplest and best systematic name of water is 'hydrogen oxide'. This is analogous to related compounds such as
hydrogen peroxide,
hydrogen sulfide, and
deuterium oxide (heavy water). Another systematic name, 'oxidane', is accepted by IUPAC as a parent name for the systematic naming of oxygen-based
substituent groups,
[14] although even these commonly have other recommended names. For example, the name
hydroxyl is recommended over ''oxidanyl'' for the –OH group. The name
oxane is explicitly mentioned by the IUPAC as being unsuitable for this purpose, since it is already the name of a cyclic ether also known as
tetrahydropyran in the
Hantzsch-Widman system; similar compounds include
dioxane and
trioxane.
Systematic nomenclature and humor
Main articles: dihydrogen monoxide hoax
Chemists sometimes refer to water as 'dihydrogen monoxide' or 'DHMO', an overly pedantic systematic covalent name of this molecule, especially in
parodies of chemical research that call for this "lethal chemical" to be banned. In reality, a more realistic systematic name would be 'hydrogen oxide', since the "di-" and "mon-" prefixes are superfluous.
Hydrogen sulfide, H
2S, is never referred to as "dihydrogen monosulfide", and
hydrogen peroxide, H
2O
2, is never called "dihydrogen dioxide".
Some overzealous
material safety data sheets for water list the following: Caution: May cause drowning!
The systematic acid name of water is 'hydroxic acid' or 'hydroxilic acid'. Likewise, the systematic alkali name of water is 'hydrogen hydroxide'—both acid and alkali names exist for water because it is able to react both as an acid or an alkali, depending on the strength of the acid or alkali it is reacted with (
amphoteric). None of these names are used widely outside of DHMO sites.
See also
★
dihydrogen monoxide hoax
★
double distilled water
★
heavy water
★
hydrodynamics
★
Mpemba effect
★
polywater theory
★
water dimer
★
Water (data page)
★
Vienna Standard Mean Ocean Water
References
1. http://www.dartmouth.edu/~etrnsfer/water.htm
2. Vienna Standard Mean Ocean Water (VSMOW), used for calibration, melts at 273.1500089(10) K (0.000089(10) °C, and boils at 373.1339 K (99.9839 °C)
3. Serway, Raymond A. ''Physics for Scientists and Engineers'', third edition
4. Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Edn.). Boca Raton (FL):CRC Press.
5. Compressibility of water as a function of temperature and pressure, Fine, R.A. and Millero, F.J., , , Journal of Chemical Physics, 1973
6. http://hyperphysics.phy-astr.gsu.edu/hbase/permot3.html
7. http://hyperphysics.phy-astr.gsu.edu/hbase/permot3.html
8. Choi 2005
9. Connolly, P.J, ''et al'', 2005
10. http://www.aip.org/enews/physnews/2003/split/648-1.html
11. http://prola.aps.org/abstract/PRL/v79/i15/p2839_1
12. http://scitation.aip.org/getabs/servlet/GetabsServlet?prog=normal&id=PRLTAO000091000005057403000001&idtype=cvips&gifs=yes
13. http://home.comcast.net/~igpl/Temperature.html
14. Leigh, G. J. ''et al.'' 1998. ''Principles of chemical nomenclature: a guide to IUPAC recommendations'', p. 99. Blackwell Science Ltd, UK. ISBN 0-86542-685-6
External links
★
Release on the IAPWS Industrial Formulation 1997 for the Thermodynamic Properties of Water and Steam (fast computation speed)
★
Release on the IAPWS Formulation 1995 for the Thermodynamic Properties of Ordinary Water Substance for General and Scientific Use (simpler formulation)
★
A spoof site on the "dangers" of dihydrogen monoxide
★
Stockholm International Water Institute (SIWI)
★
Explanation of the anomalous properties of water
★
Computational Chemistry Wiki
★
Water phase diagrams
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