Properties of water

NFPA 704
*safety square*
NFPA 704; safety square
	0 0 0

Water (H₂O)

The water molecule has this basic geometric structure

General

Systematic name

Water

Other names

Aqua
Hydrogen oxide
Hydrogen hydroxide
Hydrate
Oxidane
Hydric Acid
Dihydrogen Monoxide

Molecular formula

H₂O

Molar mass

18.0153 g·mol⁻¹

Appearance

transparent, almost
colorless liquid with
a slight hint of blue^[1]

CAS number

[7732-18-5]

Forms of water

See the Category:Forms of water

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), 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 found 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 km³ (326,000,000 mi³). Of this volume:

1 320 000 000 km³ (316,900,000 mi³ or 97.2%) is in the oceans
25 000 000 km³ (6,000,000 mi³ or 1.8%) is in glaciers and icecaps
13 000 000 km³ (3,000,000 mi³ or 0.9%) is groundwater.
250 000 km³ (60,000 mi³ or 0.02%) is fresh water in lakes, inland seas, and rivers.
13 000 km³ (3,100 mi³ 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, acetic acid, 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:

Temp (°C)	Density (g/cm³)
30	0.9957
20	0.9982
10	0.9997
0	0.9998
−10	0.9982
−20	0.9935
−30	0.9839

The values below 0 °C refer to supercooled water.

Density of saltwater and ice

The dependence of density on temperature 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 both lowers the colligative freezing point by about 2 °C and lowers the temperature of the density maximum of water to the freezing point. Hence, in ocean water, because of the salt content, the downward convection of colder water is not blocked by an expansion of water as it becomes colder near the freezing point; thus 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 bulk modulus of water is 2.2×10⁹ Pa.^[3]. The stiffness of nongasses makes them incompressible, or of less compressibility; but as their elements are not near degeneracy pressure, they are not noncompressible. Even in the deep oceans at 4000 m depth where pressures are 4×10⁷ Pa, there is only a 1.8% change in volume.^[4]

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 10⁶ volts per meter.^[5]

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.^[6]

Surface tension

Water drops are stable, due to the high surface tension of water. 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 (de-ionised) is an excellent insulator as it does not conduct electricity at all. But no water is pure due to autoionisation at any temperature above absolute zero. Because water is such a good solvent, however, it almost always has some solute dissolved in it, most frequently a salt. If water has even a tiny amount of such impurities, then it can conduct electricity, 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 H₂ gas. At the anode, four OH⁻ ions combine and release O₂ gas, molecular water, and four electrons. The gases produced bubble to the surface, where they can be collected. It is known that the theoretical maximum electrical resistivity for water is approximately 182 kilohm-m²/m (or 18.2 MΩ·cm²/cm) at 25 °C. This figure agrees well with what is typically seen on reverse osmosis, ultrafiltered and deionized ultrapure water systems used for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding that of even 100 parts per trillion (ppt) in ultrapure water will begin to noticeably lower its resistivity level by up to several kilohm-square meters/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 which makes the different charges "take out one other". 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 (H₂S), which has much weaker hydrogen bonding, is a gas at room temperature even though it has twice the molecular mass of water. The extra bonding between water molecules also gives liquid water a large specific heat capacity. This high heat capacity makes water a good heat storage medium.

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 something like 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 (H₃O⁺) 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 + H₂O ⇌ H₃O⁺ + Cl⁻

Here water is acting as a base, by receiving an H⁺ ion.

In the reaction with ammonia, NH₃, water donates an H⁺ ion, and is thus acting as an acid:

NH₃ + H₂O ⇌ NH₄⁺ + 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 CO₂ 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 H₂O at the subatomic quantum level.^[7] 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.^[8]^[9]

Water in biology

Water is essential to biology mainly due to its chemical properties on which biological processes depend on. For instance, water-soluble vitamins need water in order to be used by a body, and without water, vitamins become impotent. In addition, the solvent properties of water are vital in because many other biochemical reactions take place only within aqueous solutions (e.g., reactions in the cytoplasm and blood such as blood clotting).

History

In 1724, Gabriel Fahrenheit initially defined a temperature scale in which 60 degrees was the boiling point of water.[1]

In 1742, Anders Celsius defined the Celsius temperature scale with the freezing point of water at 100 degrees and the boiling point at standard atmospheric pressure at 0 degrees. The scale was reversed in 1744.

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,^[10] 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

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 2004, the town of Aliso Viejo, California nearly banned foam cups after learning that DHMO was used in their production.^[11] In reality, a more realistic systematic name would be hydrogen oxide, since the "di-" and "mon-" prefixes are superfluous. Hydrogen sulfide, H₂S, is never referred to as "dihydrogen monosulfide", and hydrogen peroxide, H₂O₂, 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.

There is an unconfrmed report^{[citation needed]} that in the Ottoman Empire the written formula H₂O was banned to stop people from using it as a political slogan meaning "[Sultan Abdul] Hamid II is good for nothing". That may be why some older books wrote the formula of water as OH₂.

References

^† Vienna Standard Mean Ocean Water (VSMOW), used for calibration, melts at 273.150 089(10) K (0.000 089(10) °C, and boils at 373.133 9 K (99.983 9 °C)

^ http://www.dartmouth.edu/~etrnsfer/water.htm
^ Serway, Raymond A. Physics for Scientists and Engineers, third edition
^ http://hyperphysics.phy-astr.gsu.edu/hbase/permot3.html
^ http://hyperphysics.phy-astr.gsu.edu/hbase/permot3.html
^ Choi 2005
^ Connolly, P.J, et al, 2005
^ http://www.aip.org/enews/physnews/2003/split/648-1.html
^ http://prola.aps.org/abstract/PRL/v79/i15/p2839_1
^ http://scitation.aip.org/getabs/servlet/GetabsServlet?prog=normal&id=PRLTAO000091000005057403000001&idtype=cvips&gifs=yes
^ 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
^ http://slashdot.org/articles/04/03/16/1419252.shtml?tid=133&tid=186

External links

Water Structure and Behaviour A comprehensive and up-to-date NPOV resource maintained by Prof Martin Chaplin of South Bank University, UK
A spoof site on the "dangers" of dihydrogen monoxide
Stockholm International Water Institute (SIWI)
Explanation of the anomalous properties of water
Computational Chemistry Wiki

Template:ChemicalSources

[1] ttp://www.dartmouth.edu/~etrnsfer/water.htm

[2] Serway, Raymond A. Physics for Scientists and Engineers, third edition

[3] ttp://hyperphysics.phy-astr.gsu.edu/hbase/permot3.html

[4] ttp://hyperphysics.phy-astr.gsu.edu/hbase/permot3.html

[5] Choi 2005

[6] Connolly, P.J, et al, 2005

[7] ttp://www.aip.org/enews/physnews/2003/split/648-1.html

[8] ttp://prola.aps.org/abstract/PRL/v79/i15/p2839_1

[9] ttp://scitation.aip.org/getabs/servlet/GetabsServlet?prog=normal&id=PRLTAO000091000005057403000001&idtype=cvips&gifs=yes

[10] 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

[11] ttp://slashdot.org/articles/04/03/16/1419252.shtml?tid=133&tid=186

[1]

[2]

[3]

[4]

[5]

[6]

[7]

[8]

[9]

[10]

[11]