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Properties of water

|Section2={{Chembox Properties | Formula = | MolarMass = 18.01528(33) g/mol | Odor = None | Appearance = White solid or almost colorless, transparent, with a slight hint of blue, crystalline solid or liquid | Density = Liquid: 0.9998396 g/mL at 0 °C 0.9970474 g/mL at 25 °C 0.961893 g/mL at 95 °CSolid: 0.917 g/ml ≈ 57.2 lb/ft3 | MeltingPtC = 0.00 | MeltingPt_notes = | BoilingPtC = 99.98 | BoilingPt_notes = | pKa = 13.995{{efn| A commonly quoted value of 15.7 used mainly in organic chemistry for the pKa of water is incorrect.}} | pKb = 13.995 | Viscosity = 0.890 cP | RefractIndex = 1.3330 (20°C){{Cite book|url=https://books.google.com/books?id=kTnxSi2B2FcC|title=CRC Handbook of Chemistry and Physics, 84th Edition|last=Lide|first=David R.|date=2003-06-19|publisher=CRC Press|isbn=9780849304842|series= CRC Handbook|at=8—Concentrative Properties of Aqueous Solutions: Density, Refractive Index, Freezing Point Depression, and Viscosity |language=en}} | VaporPressure = | SolubleOther = Poorly soluble in haloalkanes, aliphatic and aromatic hydrocarbons, ethers. Improved solubility in carboxylates, alcohols, ketones, amines. Miscible with methanol, ethanol, propanol, isopropanol, acetone, glycerol, 1,4-dioxane, tetrahydrofuran, sulfolane, acetaldehyde, dimethylformamide, dimethoxyethane, dimethyl sulfoxide, acetonitrile. Partially miscible with Diethyl ether, Methyl Ethyl Ketone, Dichloromethane, Ethyl Acetate, Bromine. | ThermalConductivity = 0.6065 W/m·K }} |Section3={{Chembox Structure | MolShape = Bent | CrystalStruct = Hexagonal | Dipole = 1.8546 D | PointGroup = C2v }} |Section4={{Chembox Thermochemistry | DeltaHf = -285.83 ± 0.040 kJ/mol | DeltaGf = -237.24 kJ/mol | Entropy = 69.95 ± 0.03 J/mol·K | HeatCapacity = 75.375 ± 0.05 J/mol·K }} |Section7={{Chembox Hazards | MainHazards = Drowning Avalanche (as snow)(see also Dihydrogen monoxide hoax) | NFPA-H = 0 | NFPA-F = 0 | NFPA-R = 0 | FlashPt= Non-flammable }} |Section8={{Chembox Related | OtherCations = Hydrogen sulfide Hydrogen selenide Hydrogen telluride Hydrogen polonide Hydrogen peroxide | OtherFunction_label = solvents | OtherFunction = Acetone Methanol | OtherCompounds = Water vapor Ice Heavy water }} }} Water () is a polar inorganic compound that is at room temperature a tasteless and odorless liquid, which is nearly colorless apart from a hint of blue. It is by far the most studied chemical compound and is described as the "universal solvent" for its ability to dissolve many substances. This allows it to be the "solvent of life". It is the only common substance to exist as a solid, liquid, and gas on Earth's surface. Water molecules form hydrogen bonds with each other and are strongly polar. This polarity allows it to separate ions in salts and strongly 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 °C for its molar mass, and a high heat capacity. Water is amphoteric, meaning it is both an acid and a base—it produces and more complex ions that form.}} This means the product of the concentrations of and has to be a constant. Because water is a very good solvent, it is rarely pure, and some of the properties of impure water can vary from those of the pure substance. There are also many compounds that are essentially, if not completely, insoluble in water, such as fats, oils and other non-polar substances.

Nomenclature

The accepted
IUPAC name of water is oxidane or simply water, 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 substituent nomenclature. 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. 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). The polarized form of the water molecule, , is also called hydron hydroxide by IUPAC nomenclature. Using chemical nomenclature for type I ionic binary compounds, water would take the name hydrogen monoxide, "The electropositive constituent names are cited first, and are separated from the electronegative constituent names by a space. The multiplicative prefixes may not be necessary if the oxidation states are explicit or are clearly understood." (Section 4.2.2) but this is not among the names published by the International Union of Pure and Applied Chemistry (IUPAC). It is a rarely used name of water, and mostly used in various hoaxes or spoofs that call for this "lethal chemical" to be banned. Also allowed—though the added precision is redundant—is dihydrogen monoxide such as in the dihydrogen monoxide hoax. Other systematic names for water include hydroxic acid, hydroxylic acid, and hydrogen hydroxide, using acid and base names. None of these exotic names are used widely.

Properties

Water is the chemical substance with chemical formula ; one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water vapor is essentially invisible as a gas. Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the oxygen family in the periodic table, which are gases such as hydrogen sulfide. The elements surrounding oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, all combine with hydrogen to produce gases under standard conditions. The reason that water forms a liquid is that oxygen is more electronegative than all of these elements with the exception of fluorine. Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom. These atomic charges give each water molecule a net dipole moment. Electrical attraction between water molecules due to this dipole pulls individual molecules closer together, making it more difficult to separate the molecules and therefore raising the boiling point. This attraction is known as hydrogen bonding. The molecules of water are constantly moving in relation to each other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200 femtoseconds (2×10−13 seconds). However, these bonds are strong enough to create many of the peculiar properties of water, some of which make it integral to life. Water can be described as a polar liquid that slightly dissociates disproportionately or self ionizes into an hydronium ion and hydroxide ion. 2 + The dissociation constant for this dissociation is commonly symbolized as Kw and has a value of about at 25 °C; see here for values at other temperatures.

Water, ice, and vapor

Like many substances, water can take numerous forms, which are broadly categorized by phase of matter. The liquid phase is the most common among water's phases (within the Earth's atmosphere and surface) and is the form that is generally denoted by the word "water". The solid phase of water is known as ice and commonly takes the structure of hard, amalgamated crystals, such as ice cubes, or loosely accumulated granular crystals, like snow. For a list of the many different crystalline and amorphous forms of solid , see phases of ice. The gaseous phase of water is known as water vapor (or steam), in which water takes the form of a transparent cloud. (Visible steam and clouds are, in fact, water in the liquid form as minute droplets suspended in the air.) The fourth state of water, that of a supercritical fluid, is much less common than the other three and only rarely occurs in nature, in extremely hostile conditions. When water achieves a specific critical temperature and a specific critical pressure (647 K and 22.064 MPa), the liquid and gas phases merge to one homogeneous fluid phase, with properties of both gas and liquid. 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 volcanic 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).

Heat capacity and heats of vaporization and fusion

Water has a very high specific heat capacity of 4.1814 J/(g·K) at 25 °C – the second highest among all the heteroatomic species (after ammonia), as well as a high heat of vaporization (40.65 kJ/mol or 2257 kJ/kg at the normal boiling point), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth's climate by buffering large fluctuations in temperature. According to Josh Willis, of NASA's Jet Propulsion Laboratory, the oceans can absorb one thousand times more heat than the atmosphere without changing their temperature much and are absorbing 80 to 90% of the heat of global warming. NASA – Oceans of Climate Change. Nasa.gov (2009-04-22). Retrieved on 2011-11-22. The specific enthalpy of fusion (more commonly known as latent heat) of water is 333.55 kJ/kg at 0 °C: the same amount of energy is required to melt ice as to warm ice from −160 °C up to its melting point or to heat the same amount of water by about 80 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice of glaciers and drift ice. Before and since the advent of mechanical refrigeration, ice was and still is in common use for retarding food spoilage. The specific heat capacity of ice at −10 °C is 2.03 J/(g·K) and the heat capacity of steam at 100 °C is 2.08 J/(g·K).

Density of water and ice

The density of water is about : this relationship was originally used to define the gram. The density varies with temperature, but not linearly: as the temperature increases, the density rises to a peak at and then decreases. This unusual negative thermal expansion below is also observed in molten silica. Regular, hexagonal ice is also less dense than liquid water—upon freezing, the density of water decreases by about 9%. These effects are due to the reduction of thermal motion with cooling, which allows water molecules to form more hydrogen bonds that prevent the molecules from coming close to each other. While below 4 °C the breakage of hydrogen bonds due to heating allows water molecules to pack closer despite the increase in the thermal motion (which tends to expand a liquid), above 4 °C water expands as the temperature increases. Water near the boiling point is about 4% less dense than water at . 4.135%}} Other substances that expand on freezing are silicon, gallium,Zumdahl, p. 938 germanium, bismuth, plutonium and also chemical compounds that form spacious crystal lattices with tetrahedral coordination. Under increasing pressure, ice undergoes a number of transitions to other allotropic forms with higher density than liquid water, such as ice II, ice III, high-density amorphous ice (HDA), and very-high-density amorphous ice (VHDA).The unusual density curve and lower density of ice than of water is vital to life—if water were most dense at the freezing point, then in winter the very cold water at the surface of lakes and other water bodies would sink, the lake could freeze from the bottom up, and all life in them would be killed. Furthermore, given that water is a good thermal insulator (due to its heat capacity), some frozen lakes might not completely thaw in summer. The layer of ice that floats on top insulates the water below.Zumdahl, p. 493 Water at about 4 °C (39 °F) also sinks to the bottom, thus keeping the temperature of the water at the bottom constant (see diagram).

Density of saltwater and ice

surface density]] The density of salt water 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 °C (see here for explanation) and lowers the temperature of the density maximum of water to the freezing point. This is why, in ocean water, the downward convection of colder water is not blocked by an expansion of water as it becomes colder near the freezing point. The oceans' cold water near the freezing point continues to sink. So creatures that live at the bottom of cold oceans like the Arctic Ocean generally live in water 4 °C colder than at the bottom of frozen-over fresh water lakes and rivers. 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 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 sea water just below it, in a process known as brine rejection. This denser salt water sinks by convection and the replacing seawater is subject to the same process. This produces essentially freshwater ice at −1.9 °C on the surface. The increased density of the sea water 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

Water is miscible with many liquids, for example ethanol in all proportions, forming a single homogeneous liquid. On the other hand, water and most oils are immiscible usually forming layers according to increasing density from the top. This can be predicted by comparing the 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 °C, starting at about 22 °C water will start to condense, defining the dew point, and creating fog or dew. The reverse process accounts for the fog burning off in the morning. If the humidity is increased at room temperature, for example, by running a hot shower or a bath, and the temperature stays about the same, the vapor soon reaches the pressure for phase change, and then condenses out as minute water droplets, commonly referred to as steam. A gas in this context is referred to as saturated or 100% relative humidity, when the vapor pressure of water in the air is at the equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool enough) will fail to lose mass through evaporation when exposed to saturated air. Because the amount of water vapor in air is small, relative humidity, the ratio of the partial pressure due to the water vapor to the saturated partial vapor pressure, is much more useful. Water vapor pressure above 100% relative humidity is called super-saturated and can occur if air is rapidly cooled, for example, by rising suddenly in an updraft.

Vapor pressure

Compressibility

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

Triple point

3D representation of several phases of water.jpg|A 3-dimensional representation of the relationships between allotropes of ice amongst each other and in relation to liquid water and gaseous water. The temperature and pressure at which solid, liquid, and gaseous water coexist in equilibrium is called the triple point of water. This point is used to define the units of temperature (the kelvin, the SI unit of thermodynamic temperature and, indirectly, the degree Celsius and even the degree Fahrenheit). As a consequence, water's triple point temperature, as measured in these units, is a prescribed value rather than a measured quantity. This pressure is quite low, about of the normal sea level barometric pressure of 101,325 Pa. The atmospheric surface pressure on planet Mars is 610.5 Pa, which is remarkably close to the triple point pressure. The altitude of this surface pressure was used to define zero-elevation or "sea level" on that planet. Although it is commonly named as "the triple point of water", the stable combination of liquid water, ice I, and water vapor is but one of several triple points on the phase diagram of water. Gustav Heinrich Johann Apollon Tammann in Göttingen produced data on several other triple points in the early 20th century. Kamb and others documented further triple points in the 1960s.

Melting point

The melting point of ice is 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 . The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, by /atm{{efn| name = sharp | The source gives it as 0.0072°C/atm. However the author defines an atmosphere as 1,000,000 dynes/cm2 (a bar). Using the standard definition of atmosphere, 1,013,250 dynes/cm2, it works out to 0.0073°C/atm}} or about /70 atm{{efn|name = sharp2| Using the fact that 0.5/0.0073 68.5}} as the stabilization energy of hydrogen bonding is exceeded by intermolecular repulsion, but as ice transforms into its allotropes (see crystalline states of ice) above , the melting point increases markedly with pressure, i.e., reaching at (triple point of Ice VII).

Electrical properties

Electrical conductivity

Pure water containing no exogenous ions is an excellent insulator, but not even "deionized" water is completely free of ions. Water undergoes auto-ionization in the liquid state, when two water molecules form one hydroxide anion () and one hydronium cation (). Because water is such a good solvent, it almost always has some solute dissolved in it, often a salt. If water has even a tiny amount of such an impurity, then it can conduct electricity far more readily. It is known that the theoretical maximum electrical resistivity for water is approximately 18.2 MΩ·cm (182 ·m) at 25 °C. This figure agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in otherwise ultra-pure water begins to noticeably lower its resistivity by up to several kΩ·m. In pure water, sensitive equipment can detect a very slight electrical conductivity of 0.05501 ± 0.0001 µS/ cm at 25.00 °C. Water can also be electrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is a very slow process, as very little current is conducted. In ice, the primary charge carriers are protons (see proton conductor).

Polarity, hydrogen bonding and intermolecular structure

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°, which is smaller than the typical tetrahedral angle of 109.47°. The lone pairs are closer to the oxygen atom than the electrons sigma bonded to the hydrogens, so they require more space. The increased repulsion of the lone pairs forces the O–H bonds closer to each other.Zumdahl, p. 393 Another effect of the electronic structure is that water is a polar molecule. Due to the difference in electronegativity, there is a bond dipole moment pointing from each H to the O, making the oxygen partially negative and each hydrogen partially positive. In addition, the lone pairs of electrons on the O are in the direction opposite to the hydrogen atoms. This results in a large molecular dipole, pointing from a positive region between the two hydrogen atoms to the negative region of the oxygen atom. The charge differences cause water molecules to be attracted to each other (the relatively positive areas being attracted to the relatively negative areas) and to other polar molecules. This attraction contributes to hydrogen bonding, and explains many of the properties of water, such as solvent action. 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. These properties include its relatively high melting and boiling point temperatures: more energy is required to break the hydrogen bonds between water molecules. In contrast, hydrogen sulfide (), has much weaker hydrogen bonding due to sulfur's lower electronegativity. is a gas at room temperature, in spite of 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.

Proposed structures

s (1) between molecules of water]] A single water molecule can participate in a maximum of four hydrogen bonds 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 thermodynamic, 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 °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 liquid form typically bind not to four but to 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. Water, the team suggests, is a muddle of the two proposed structures. They say that it is a soup flecked with "icebergs" each comprising 100 or so loosely connected molecules that are relatively open and hydrogen bonded. The soup is made of the string structure and the icebergs of the tetrahedral structure.

Cohesion and adhesion

drops adhering to a spider web]] Water molecules stay close to each other ( 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. Water also has high adhesion properties because of its polar nature. On extremely 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. In biological cells and organelles, 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. They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.

Surface tension

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.]] Water has a high surface tension of 71.99 mN/m at 25 °C, | 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 | accessdate = 7 Jan 2011 }} or sweet flavor.{{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 | accessdate = 7 Jan 2011}} 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. Oxygen also has three stable isotopes, with present in 99.76%, in 0.04%, and in 0.2% of water molecules. Light water refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard level.

Standard 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 protium. Only 155 ppm include deuterium ( or D), a hydrogen isotope with one neutron, and fewer than 20 parts per quintillion include tritium ( or T), which has two neutrons.

Reactions

Acid-base reactions

Water is amphoteric: it has the ability to act as either an acid or a base in chemical reactions.Zumdahl, p. 659 According to the Brønsted-Lowry definition, an acid is a proton () donor and a base is a proton acceptor.Zumdahl, p. 654 When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid. For instance, water receives an {{chem|link= |H|+}} ion from HCl when hydrochloric acid is formed: + }} + In the reaction with ammonia, , water donates a ion, and is thus acting as an acid: }} + }} + Because the oxygen atom in water has two lone pairs, water often acts as a Lewis base, or electron pair donor, in reactions with Lewis acids, although it can also react with Lewis bases, forming hydrogen bonds between the electron pair donors and the hydrogen atoms of water. HSAB theory describes water as both a weak hard acid and a weak hard base, meaning that it reacts preferentially with other hard species: }} + }} → }} + }} → )|6|3+}} }} + }} → )|6|-}} When a salt of a weak acid or of a weak base is dissolved in water, water can partially hydrolyze the salt, producing the corresponding base or acid, which gives aqueous solutions of soap and baking soda their basic pH: + NaOH +

Ligand chemistry

Water's Lewis base character makes it a common ligand in transition metal complexes, examples of which range from solvated ions, such as , to perrhenic acid, which contains two water molecules coordinated to a rhenium atom, and various solid hydrates, such as . Water is typically a monodentate ligand–it forms only one bond with the central atom.Zumdahl, p. 984

Organic chemistry

As a hard base, water reacts readily with organic carbocations; for example in an hydration reaction, a hydroxyl group () and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an alcohol. When 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 substitution and 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.Zumdahl, p. 171 It oxidizes chemicals such as hydrides, alkali metals, and some alkaline earth metals.Zumdahl, pp. 932, 936 One example of an alkali metal reacting with water is:Zumdahl, p. 338 2 Na + 2 → + 2 + 2 Some other reactive metals, such as aluminum and beryllium, are oxidized by water as well, but their oxides adhere to the metal and form a passive protective layer.Zumdahl, p. 862 Note, however, that the rusting of iron is a reaction between iron and oxygenZumdahl, p. 981 that is dissolved in water, not between iron and water. 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 . Almost all such reactions require a catalyst. An example of the oxidation of water is: 4 + 2 → 4 AgF + 4 HF +

Electrolysis

Water can be split into its constituent elements, hydrogen and oxygen, by passing an electric current through it. This process is called electrolysis. The cathode half reaction is: 2 + 2 → The anode half reaction is: 2 → + 4 + 4 The gases produced bubble to the surface, where they can be collected. The standard potential of the water electrolysis cell (when heat is added to the reaction) is a minimum of 1.23 V at 25 °C. The operating potential is actually 1.48 V (or above) in practical electrolysis when heat input is negligible.

History

Henry Cavendish showed that water was composed of oxygen and hydrogen in 1781. The first decomposition of water into hydrogen and oxygen, by electrolysis, was done in 1800 by English chemist William Nicholson and Anthony Carlisle. In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen. Gilbert Newton Lewis isolated the first sample of pure heavy water in 1933. The properties of water have historically been used to define various temperature scales. Notably, the Kelvin, Celsius, Rankine, 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. A Brief History of Temperature Measurement. Home.comcast.net. Retrieved on 2011-11-22.

See also

Notes

References

Cited sources

External links

"green air" © 2007 - Ingo Malchow, Webdesign Neustrelitz
This article based upon the http://en.wikipedia.org/wiki/Properties_of_water, the free encyclopaedia Wikipedia and is licensed under the GNU Free Documentation License.
Further informations available on the list of authors and history: http://en.wikipedia.org/w/index.php?title=Properties_of_water&action=history
presented by: Ingo Malchow, Mirower Bogen 22, 17235 Neustrelitz, Germany